© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
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CORROSION UNDERSTANDING THE BASICS
Edited by J.R. Davis Davis & Associates
ASM International® Materials Park, Ohio 44073-0002
© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
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Copyright Ó 2000 by ASM International® All rights reserved No part of this book may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise, without the written permission of the copyright owner. First printing, January 2000 Great care is taken in the compilation and production of this book, but it should be made clear that NO WARRANTIES, EXPRESS OR IMPLIED, INCLUDING, WITHOUT LIMITATION, WARRANTIES OF MERCHANTABILITY OR FITNESS FOR A PARTICULAR PURPOSE, ARE GIVEN IN CONNECTION WITH THIS PUBLICATION. Although this information is believed to be accurate by ASM, ASM cannot guarantee that favorable results will be obtained from the use of this publication alone. This publication is intended for use by persons having technical skill, at their sole discretion and risk. Since the conditions of product or material use are outside of ASM’s control, ASM assumes no liability or obligation in connection with any use of this information. No claim of any kind, whether as to products or information in this publication, and whether or not based on negligence, shall be greater in amount than the purchase price of this product or publication in respect of which damages are claimed. THE REMEDY HEREBY PROVIDED SHALL BE THE EXCLUSIVE AND SOLE REMEDY OF BUYER, AND IN NO EVENT SHALL EITHER PARTY BE LIABLE FOR SPECIAL, INDIRECT OR CONSEQUENTIAL DAMAGES WHETHER OR NOT CAUSED BY OR RESULTING FROM THE NEGLIGENCE OF SUCH PARTY. As with any material, evaluation of the material under end-use conditions prior to specification is essential. Therefore, specific testing under actual conditions is recommended. Nothing contained in this book shall be construed as a grant of any right of manufacture, sale, use, or reproduction, in connection with any method, process, apparatus, product, composition, or system, whether or not covered by letters patent, copyright, or trademark, and nothing contained in this book shall be construed as a defense against any alleged infringement of letters patent, copyright, or trademark, or as a defense against liability for such infringement. Comments, criticisms, and suggestions are invited, and should be forwarded to ASM International. ASM International staff who worked on this project included Scott Henry, Assistant Director, Reference Publications; Bonnie Sanders, Manager of Copy Editing; Grace Davidson, Manager of Book Production; Nancy Hrivnak and Carol Terman, Copy Editors; Candace Mullet and Jill Kinson, Book Production Coordinators. Library of Congress Cataloging-in-Publication Data Corrosion: understanding the basics / edited by J.R. Davis. p. cm. Includes bibliographical references and index. 1. Corrosion and anti-corrosives. I. Davis, J.R. (Joseph R.) TA462.C668 2000 620.1’1223—dc21 99-057146 ISBN: 0-87170-641-5 SAN: 204-7586 ASM International® Materials Park, OH 44073-0002 http://www.asm-intl.org Printed in the United States of America
© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
Contents Preface ...................................................................................ix CHAPTER 1: The Effects and Economic Impact of Corrosion .....1 The Definition of Corrosion.............................................................2 The Effects of Corrosion ..................................................................3 The Many Forms of Corrosion.........................................................4 Methods to Control Corrosion .........................................................6 Material Selection...........................................................................6 Coatings...........................................................................................7 Inhibitors .........................................................................................8 Cathodic Protection ........................................................................8 Design..............................................................................................8 Opportunities in Corrosion Control.................................................9 The Economic Impact of Corrosion ..............................................10 Sources of Information ...................................................................14 Appendix: Addresses of Trade Associations and Technical Societies Involved with Corrosion ..........................17 CHAPTER 2: Basic Concepts Important to Corrosion .........21 Behavior of a Metal in an Environment ........................................21 The Four Requirements of a Corrosion Cell.................................23 Metal Characteristics Important to Corrosion ..............................25 Metallurgical Characteristics .......................................................25 Inherent Reactivity .......................................................................35 Formation of Corrosion Products ................................................37 Important Solution Characteristics ................................................38 Corrosion Rate Expressions and Allowances ...............................45 CHAPTER 3: Principles of Aqueous Corrosion ....................49 The Thermodynamics of Aqueous Corrosion ...............................50 iii
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Corrosion Reactions and Free-Energy Change...........................50 Free Energy and Electrochemical Potential................................53 Tendency for Metals to Corrode..................................................55 Effect of Ionic Concentration on Electrode Potential ................56 Electromotive Force Series ..........................................................59 Galvanic Series .............................................................................60 Standard Electrode Potentials for Other Reactions ....................62 Potential-pH Diagrams: General Aspects ...................................62 Potential-pH Diagrams for Specific Metals................................67 Strategies for Corrosion Control from E-pH Diagrams .............74 Limitations of E-pH Diagrams ....................................................76 The Kinetics of Aqueous Corrosion ..............................................77 Electrochemical Reactions...........................................................77 Mixed-Potential Theory ...............................................................79 Types of Polarization ...................................................................82 Applications of Mixed-Potential Theory Diagrams ...................88 Exchange Currents........................................................................95 CHAPTER 4: Forms of Corrosion: Recognition and Prevention.....99 Uniform Corrosion .......................................................................100 Pitting Corrosion...........................................................................102 Crevice Corrosion.........................................................................107 Tuberculation ..............................................................................114 Deposit Corrosion.......................................................................118 Filiform Corrosion......................................................................122 Poultice Corrosion ......................................................................125 Galvanic Corrosion.......................................................................125 General Description....................................................................125 Galvanic Series ...........................................................................126 Polarization .................................................................................129 Factors Influencing Galvanic Corrosion Behavior...................129 Situations That Promote Galvanic Attack ................................130 Prevention of Galvanic Corrosion .............................................133 Erosion-Corrosion ........................................................................134 General Description....................................................................134 Critical Factors Influencing Erosion-Corrosion .......................137 Prevention of Erosion-Corrosion...............................................144 Cavitation ....................................................................................146 Fretting Corrosion ......................................................................149 Intergranular Corrosion ................................................................151 General Description....................................................................151 Intergranular Corrosion of Austenitic Stainless Steels ............152 Intergranular Corrosion of Other Alloy Systems .....................155 Exfoliation ..................................................................................157 Dealloying Corrosion ...................................................................158 Dezincification............................................................................158 Graphitic Corrosion....................................................................162 Stress-Corrosion Cracking ...........................................................164 Corrosion Fatigue .........................................................................175
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© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
Hydrogen Damage ........................................................................180 Hydrogen Embrittlement............................................................180 Hydrogen-Induced Blistering ....................................................184 Cracking from Precipitation of Internal Hydrogen ..................185 Hydrogen Attack.........................................................................186 Hydride Formation .....................................................................187 Prevention of Hydrogen Damage ..............................................188 Liquid-Metal Embrittlement ........................................................189 CHAPTER 5: Types of Corrosive Environments .................193 Characteristics of Corrosive Environments ................................194 Biologically Influenced Corrosion ..............................................199 Industries and Organisms Involved...........................................200 Tuberculation ..............................................................................203 Prevention of MIC ......................................................................204 Atmospheric Corrosion ................................................................205 Underground/Soil Corrosion........................................................211 Factors Affecting Underground/Soil Corrosion .......................211 Types of Underground/Soil Corrosion......................................213 Corrosion Control .......................................................................215 Natural and Treated Waters .........................................................216 Understanding Corrosion in Acids ..............................................217 Corrosion by Sulfuric Acid ..........................................................220 Materials Selection Guidelines for Sulfuric Acid ....................220 Use of Steel in Sulfuric Acid .....................................................221 Use of Cast Irons in Sulfuric Acid ............................................223 Use of Stainless Steels in Sulfuric Acid ...................................223 Use of Nickel Alloys in Sulfuric Acid ......................................224 Other Metals Used in Sulfuric Acid ..........................................225 Nonmetallic Materials Used in Sulfuric Acid ..........................225 Corrosion by Nitric Acid..............................................................226 Materials Selection Guidelines for Nitric Acid ........................227 Corrosion by Hydrochloric Acid .................................................227 Materials Selection Guidelines for Hydrochloric Acid ...........228 Corrosion by Hydrogen Fluoride and Hydrofluoric Acid..........228 Materials Selection Guidelines for Hydrofluoric Acid ............229 Corrosion by Phosphoric Acid.....................................................230 Materials Selection Guidelines for Phosphoric Acid ...............231 Corrosion by Organic Acids ........................................................231 Acetic Acid .................................................................................232 Other Organic Acids...................................................................234 Corrosion by Alkalis ....................................................................234 Materials Selection Guidelines for Alkalis...............................234 CHAPTER 6: Corrosion Characteristics of Structural Materials.....................................................237 Carbon Steels ................................................................................238 Corrosive Service .......................................................................238 Protection of Steel from Corrosion ...........................................239 v
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Weathering Steels .........................................................................242 Alloy Steels ...................................................................................244 Cast Irons ......................................................................................244 Commercially Available Cast Irons ..........................................245 Graphitic Corrosion....................................................................246 Stainless Steels..............................................................................247 Stainless Steel Families..............................................................247 Mechanism of Corrosion Resistance.........................................252 Forms of Corrosion of Stainless Steels .....................................253 Corrosion in Various Applications............................................256 Nickel and Nickel-Base Alloys ...................................................259 Effects of Major Alloying Elements .........................................260 Chemical-Processing Applications............................................262 Seawater Applications................................................................263 Applications in Pulp and Paper Mills .......................................264 Flue Gas Desulfurization Applications .....................................265 Sour Gas Applications................................................................265 High-Temperature Applications ................................................265 Copper and Copper-Base Alloys .................................................266 Effects of Alloy Composition....................................................267 Types of Attack ..........................................................................269 Applications of Copper-Base Alloys.........................................269 Aluminum and Aluminum-Base Alloys ......................................270 Effects of Alloy Composition....................................................271 Modes of Corrosion That Attack Aluminum ............................272 Corrosion Protection of Aluminum ...........................................275 Applications of Aluminum-Base Alloys...................................277 Titanium and Titanium-Base Alloys ...........................................278 Mechanism of Corrosion Resistance.........................................279 Modes of Corrosion That Attack Titanium...............................280 Corrosion Protection of Titanium..............................................281 Applications of Titanium-Base Alloys .....................................281 Zinc and Zinc-Base Alloys ..........................................................282 Magnesium and Magnesium-Base Alloys...................................282 Lead and Lead Alloys...................................................................284 Tin and Tin-Base Alloys ..............................................................286 Zirconium and Zirconium-Base Alloys.......................................287 Tantalum........................................................................................287 Niobium and Niobium-Base Alloys ............................................288 Cobalt-Base Alloys.......................................................................289 Polymers........................................................................................289 Types of Polymers ......................................................................290 Properties of Polymers ...............................................................290 Environmental Degradation of Polymers..................................291 Ceramics........................................................................................295 Other Nonmetallic Materials........................................................297 Rubber .........................................................................................297 Carbon and Graphite ..................................................................299 Woods..........................................................................................299
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© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
CHAPTER 7: Corrosion Control by Proper Design ............301 Design as a Process ......................................................................302 The Design Team........................................................................302 Steps in the Design Process .......................................................303 General Considerations in Corrosion-Control Design ...............303 Design Details that Accelerate Corrosion...................................308 Design Solutions for Specific Forms of Corrosion ....................320 Corrosion Allowance....................................................................324 Design Considerations for Using Weathering Steels .................325 Failures Involving Corrosion of Structural Steel .....................326 CHAPTER 8: Corrosion Control by Materials Selection ....331 Elements of the Materials Selection Process ..............................333 Materials Considerations..............................................................341 Selecting Materials to Avoid or Minimize Corrosion ................349 General Corrosion ......................................................................353 Localized Corrosion ...................................................................358 CHAPTER 9: Corrosion Control by Protective Coatings and Inhibitors ..............................................................363 Organic Coatings and Linings .....................................................364 Design and Selection of a Coating System ...............................365 Surface Preparation ....................................................................367 Inspection and Quality Assurance .............................................369 Coating and Lining Materials ....................................................371 Environmental, Health, and Safety Considerations .................379 Metallic Coatings..........................................................................382 Electroplated Coatings ...............................................................382 Electroless Nickel Plating .........................................................386 Hot-Dip Coatings........................................................................387 Thermal Spray Coatings.............................................................391 Clad Metals .................................................................................392 Pack Cementation .......................................................................394 Vapor-Deposited Coatings.........................................................395 Surface Modification..................................................................395 Nonmetallic Inorganic Coatings ..................................................396 Concrete and Cementatious Coatings and Linings...................397 Porcelain Enamels ......................................................................398 Conversion Coatings ..................................................................399 Aluminum Anodizing.................................................................401 Inhibitors .......................................................................................401 Types of Inhibitors .....................................................................402 Biocides.......................................................................................404 Application of Inhibitors............................................................405 CHAPTER 10: Corrosion Control by Cathodic and Anodic Protection ................................................407 Cathodic Protection ......................................................................407 How Cathodic Protection Works ...............................................408 vii
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Types of Cathodic Protection ....................................................410 Anode Materials .........................................................................411 Criteria for Cathodic Protection ................................................414 Problems with Cathodic Protection...........................................415 Applications of Cathodic Protection .........................................417 Anodic Protection .........................................................................422 The Concept of Anodic Protection ............................................422 Equipment Required for Anodic Protection .............................423 Applications of Anodic Protection ............................................425 CHAPTER 11: Corrosion Testing and Monitoring..............427 Classification of Corrosion Testing.............................................427 Purposes of Corrosion Tests ........................................................429 Steps in a Corrosion Test Program ..............................................430 Preparation and Cleaning of Test Specimens .............................432 Specific Types of Laboratory Tests.............................................433 Simulated Atmosphere Tests .....................................................434 Salt-Spray Testing ......................................................................435 Immersion Tests .........................................................................438 Field Tests .....................................................................................441 Atmospheric Tests ......................................................................442 Electrochemical Tests...................................................................448 Electrochemical Test Classification ..........................................448 Reference Electrode ...................................................................449 Types of Electrochemical Measurements .................................451 Applications of Electrochemical Tests .....................................456 Corrosion Monitoring...................................................................467 Selecting a Corrosion-Monitoring Method...............................470 Strategies in Corrosion Monitoring...........................................472 CHAPTER 12: Techniques for Diagnosis of Corrosion Failures .......................................................475 Factors That Influence Corrosion Failures .................................475 Analysis of Corrosion Failures ....................................................481 Collection of Background Data .................................................482 On-Site Examination ..................................................................483 On-Site Sampling .......................................................................483 Preliminary Laboratory Examination........................................484 Microscopic Examination ..........................................................485 Chemical Analysis......................................................................486 Bulk Material Analysis ..............................................................488 Nondestructive Evaluation.........................................................489 Corrosion Testing .......................................................................490 Mechanical Testing ....................................................................491 Analyzing the Evidence, Formulating Conclusions, and Writing the Report...........................................................492 APPENDIX 1: Glossary of Corrosion-Related Terms .........497 Index ...................................................................................517 viii
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Preface Most people are familiar with corrosion in some form or another. Whether it is a rusty nail in a backyard fence, corroded fenders and/or mufflers on our automobiles, or a perforated underground water pipe, it is safe to say that corrosion is all around us. It is costly to prevent or repair, and it is generally not pleasing to look at. In the industrial workplace, corrosion is certainly one of the most common causes of failure of engineered components and structures. The complexities of corrosion phenomena challenge corrosion scientists, chemists, mechanical, civil, and metallurgical engineers, coating specialists, and maintenance and operating personnel. In order to better understand corrosion, it is important to first examine the basic concepts that influence the corrosion process; hence, the title of this publication—Corrosion: Understanding the Basics. Included in these 12 chapters are practical discussions on the following: · Thermodynamic and electrochemical principles of corrosion · Recognition and prevention of various forms of corrosion · Types of corrosive environments commonly encountered and environmental variables that can increase or decrease corrosion rates · Corrosion characteristics of metals and alloys and nonmetallic materials · Methods of corrosion prevention, including design considerations, materials selection, coatings, inhibitors, and cathodic and anodic protection · Corrosion testing and monitoring · Techniques for diagnosing corrosion failures
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© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
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Although the book is primarily intended for professionals who are not corrosion experts, it should also serve as a quick and useful corrosioncontrol guide for corrosion engineers. Assisting in the preparation of this book was Larry Korb from Rockwell International. Larry, who is a Fellow of ASM International and longtime member and former chairman of the ASM Handbook Committee, meticulously reviewed each chapter. I have long been in awe of my friend’s exhaustive knowledge of materials and their failure mechanisms (including corrosion), and his keen insight into the editorial process. It is always an honor and a privilege to work with Mr. Korb. I also wish to acknowledge the contributions of Nalco Chemical Company (Naperville, IL). Many of the photographs illustrating the different modes of corrosion were supplied by Nalco. These originally appeared in two excellent books on failure analysis authored by Nalco engineers Harvey M. Herro (an ASM member) and Robert D. Port. I am indebted to Ms. Connie Szewczyk, a Communications Specialists with Nalco, for supplying these photographs. Thanks are also extended to Kenneth B. Tator and Alison B. Kaelin from KTA-Tator Inc. (Pittsburgh, PA). Ken supplied an extensive table that reviewed the advantages and limitations of organic coating resins. Alison prepared material on environmental, health, and safety considerations for the coatings industry. Their contributions appear in Chapter 9. The efforts of the ASM staff are also duly noted. In particular, I would like to thank Scott Henry and Bonnie Sanders from the Publications Department and Eleanor Baldwin and her coworkers from the ASM Library for the help and support throughout the project Last, I would be remiss in not acknowledging the fact that several chapters in the book were adapted from the ASM Materials Engineering Institute (MEI) course on corrosion that was prepared by Dr. Joe H. Payer from Case Western Reserve University (Cleveland, OH). Chapters 2 and 3, as well as the description of electrochemical test methods in Chapter 11, were based on Dr. Payer’s work. Joseph R. Davis Davis & Associates Chagrin Falls, Ohio
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© 2000 ASM International. All Rights Reserved. Corrosion: Understanding the Basics (#06691G)
CHAPTER
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1
The Effects and Economic Impact of Corrosion CORROSION is a natural process. Just like water flows to the lowest level, all natural processes tend toward the lowest possible energy states. Thus, for example, iron and steel have a natural tendency to combine with other chemical elements to return to their lowest energy states. In order to return to lower energy states, iron and steel frequently combine with oxygen and water, both of which are present in most natural environments, to form hydrated iron oxides (rust), similar in chemical composition to the original iron ore. Figure 1 illustrates the corrosion life cycle of a steel product. Finished Steel Product
Smelting & Refining
Air & Moisture Corrode Steel & Form Rust
Adding Energy
Giving Up Energy
Mining Ore
Iron Oxide (Ore & Rust)
Fig. 1
The corrosion cycle of steel
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The Definition of Corrosion Corrosion can be defined in many ways. Some definitions are very narrow and deal with a specific form of corrosion, while others are quite broad and cover many forms of deterioration. The word corrode is derived from the Latin corrodere, which means “to gnaw to pieces.” The general definition of corrode is to eat into or wear away gradually, as if by gnawing. For purposes here, corrosion can be defined as a chemical or electrochemical reaction between a material, usually a metal, and its environment that produces a deterioration of the material and its properties. The environment consists of the entire surrounding in contact with the material. The primary factors to describe the environment are the following: (a) physical state—gas, liquid, or solid; (b) chemical composition— constituents and concentrations; and (c) temperature. Other factors can be important in specific cases. Examples of these factors are the relative velocity of a solution (because of flow or agitation) and mechanical loads on the material, including residual stress within the material. The emphasis in this chapter, as well as in other chapters in this book, is on aqueous corrosion, or corrosion in environments where water is present. The deterioration of materials because of a reaction with hot gases, however, is included in the definition of corrosion given here. To summarize, corrosion is the deterioration of a metal and is caused by the reaction of the metal with the environment. Reference to marine corrosion of a pier piling means that the steel piling corrodes because of its reaction with the marine environment. The environment is airsaturated seawater. The environment can be further described by specifying the chemical analysis of the seawater and the temperature and velocity of the seawater at the piling surface. When corrosion is discussed, it is important to think of a combination of a material and an environment. The corrosion behavior of a material cannot be described unless the environment in which the material is to be exposed is identified. Similarly, the corrosivity or aggressiveness of an environment cannot be described unless the material that is to be exposed to that environment is identified. In summary, the corrosion behavior of the material depends on the environment to which it is subjected, and the corrosivity of an environment depends on the material exposed to that environment. It is useful to identify both natural combinations and unnatural combinations in corrosion. Examples of natural or desirable combinations of material and environment include nickel in caustic environments, lead in water, and aluminum in atmospheric exposures. In these environments, the interaction between the metal and the environment does not
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The Effects and Economic Impact of Corrosion
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usually result in detrimental or costly corrosion problems. The combination is a natural combination to provide good corrosion service. Unnatural combinations, on the other hand, are those that result in severe corrosion damage to the metal because of exposure to an undesirable environment. Examples of unnatural combinations include copper in ammonia solutions, stainless steel in chloride-containing environments (e.g., seawater), and lead with wine (acetic acid in wine attacks lead). It has been postulated that the downfall of the Roman Empire can be attributed in part to a corrosion problem, specifically the storage of wine in lead-lined vessels. Lead dissolved in the wine and consumed by the Roman hierarchy resulted in insanity (lead poisoning) and contributed to the subsequent eventual downfall. Another anecdote regarding lead and alcoholic beverages dates back to the era of Benjamin Franklin. One manifestation was the “dry bellyache” with accompanying paralysis, which was mentioned by Franklin in a letter to a friend. This malady was actually caused by the ingestion of lead from corroded lead coil condensers used in making brandy. The problem became so widespread that the Massachusetts legislature passed a law in the late 1700s that outlawed the use of lead in producing alcoholic beverages.
The Effects of Corrosion
The effects of corrosion in our daily lives are both direct, in that corrosion affects the useful service lives of our possessions, and indirect, in that producers and suppliers of goods and services incur corrosion costs, which they pass on to consumers. At home, corrosion is readily recognized on automobile body panels, charcoal grills, outdoor furniture, and metal tools. Preventative maintenance such as painting protects such items from corrosion. A principal reason to replace automobile radiator coolant every 12 to 18 months is to replenish the corrosion inhibitor that controls corrosion of the cooling system. Corrosion protection is built into all major household appliances such as water heaters, furnaces, ranges, washers, and dryers. Of far more serious consequence is how corrosion affects our lives during travel from home to work or school. The corrosion of steel reinforcing bar (rebar) in concrete can proceed out of sight and suddenly (or seemingly so) result in failure of a section of highway, the collapse of electrical towers, and damage to buildings, parking structures, and bridges, etc., resulting in significant repair costs and endangering public safety. For example, the sudden collapse because of corrosion fatigue of the Silver Bridge over the Ohio River at Point Pleasant, OH in 1967 resulted in the loss of 46 lives and cost millions of dollars.
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Perhaps most dangerous of all is corrosion that occurs in major industrial plants, such as electrical power plants or chemical processing plants. Plant shutdowns can and do occur as a result of corrosion. This is just one of its many direct and indirect consequences. Some consequences are economic, and cause the following: · · · · · ·
Replacement of corroded equipment Overdesign to allow for corrosion Preventive maintenance, for example, painting Shutdown of equipment due to corrosion failure Contamination of a product Loss of efficiency—such as when overdesign and corrosion products decrease the heat-transfer rate in heat exchangers · Loss of valuable product, for example, from a container that has corroded through · Inability to use otherwise desirable materials · Damage of equipment adjacent to that in which corrosion failure occurs Still other consequences are social. These can involve the following issues: · Safety, for example, sudden failure can cause fire, explosion, release of toxic product, and construction collapse · Health, for example, pollution due to escaping product from corroded equipment or due to a corrosion product itself · Depletion of natural resources, including metals and the fuels used to manufacture them · Appearance as when corroded material is unpleasing to the eye
Of course, all the preceding social items have economic aspects also (see the discussion that follows, “Economic Impact of Corrosion”). Clearly, there are many reasons for wanting to avoid corrosion.
The Many Forms of Corrosion Corrosion occurs in several widely differing forms. Classification is usually based on one of three factors: · Nature of the corrodent: Corrosion can be classified as “wet” or “dry.” A liquid or moisture is necessary for the former, and dry corrosion usually involves reaction with high-temperature gases. · Mechanism of corrosion: This involves either electrochemical or direct chemical reactions.
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· Appearance of the corroded metal: Corrosion is either uniform and the metal corrodes at the same rate over the entire surface, or it is localized, in which case only small areas are affected.
Classification by appearance, which is particularly useful in failure analysis, is based on identifying forms of corrosion by visual observation with either the naked eye or magnification. The morphology of attack is the basis for classification. Figure 2 illustrates schematically some of the most common forms of corrosion. Eight forms of wet (or aqueous) corrosion can be identified based on appearance of the corroded metal. These are: · Uniform or general corrosion · Pitting corrosion · Crevice corrosion, including corrosion under tubercles or deposits, filiform corrosion, and poultice corrosion · Galvanic corrosion · Erosion-corrosion, including cavitation erosion and fretting corrosion · Intergranular corrosion, including sensitization and exfoliation · Dealloying, including dezincification and graphitic corrosion · Environmentally assisted cracking, including stress-corrosion cracking, corrosion fatigue, and hydrogen damage
In theory, the eight forms of corrosion are clearly distinct; in practice however, there are corrosion cases that fit in more than one category. Other corrosion cases do not appear to fit well in any of the eight categories. Nevertheless, this classification system is quite helpful in the study Load More noble metal
No corrosion
Uniform
Galvanic
Flowing corrodent
Cyclic movement
Erosion
Fretting
Tensile stress
Pitting
Fig. 2
Exfoliation
Dealloying
Intergranular
Stress-corrosion cracking
Schematics of the common forms of corrosion
Metal or nonmetal
Crevice Cyclic stress
Corrosion fatigue
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Corrosion: Understanding the Basics
CORROSION
UNIFORM
LOCALIZED
MACROSCOPIC Galvanic
MICROSCOPIC
Erosion-corrosion Crevice
Intergranular
Pitting
Stress-corrosion cracking
Exfoliation
Corrosion fatigue
Dealloying
Fig. 3
Macroscopic versus microscopic forms of localized corrosion
of corrosion problems. Detailed information on these eight forms of corrosion can be found in Chapter 4. Completeness requires further distinction between macroscopically localized corrosion and microscopic local attack. In the latter case, the amount of metal dissolved is minute, and considerable damage can occur before the problem becomes visible to the naked eye. Macroscopic forms of corrosion affect greater areas of corroded metal and are generally observable with the naked eye or can be viewed with the aid of a low-power magnifying device. Figure 3 classifies macroscopic and microscopic forms of localized corrosion.
Methods to Control Corrosion There are five primary methods of corrosion control: · · · · ·
Material selection Coatings Inhibitors Cathodic protection Design
Each is described briefly here and in more detail in subsequent chapters.
Material Selection Each metal and alloy has unique and inherent corrosion behavior that can range from the high resistance of noble metals, for example, gold
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The Effects and Economic Impact of Corrosion
and platinum, to the low corrosion resistance of active metals, for example, sodium and magnesium. Furthermore, the corrosion resistance of a metal strongly depends on the environment to which it is exposed, that is, the chemical composition, temperature, velocity, and so forth. The general relation between the rate of corrosion, the corrosivity of the environment, and the corrosion resistance of a material is: corrosivity of environment » rate of corrosive attack corrosion resistance of metal For a given corrosion resistance of the material, as the corrosivity of the environment increases, the rate of corrosion increases. For a given corrosivity of the environment, as the corrosion resistance of the material increases, the rate of corrosion decreases. Often an acceptable rate of corrosion is fixed and the challenge is to match the corrosion resistance of the material and the corrosivity of the environment to be at or below the specified corrosion rate. Often there are several competing materials that can meet the corrosion requirements, and the material selection process becomes one of determining which of the candidate materials provides the most economical solution for the particular service. Consideration of corrosion resistance is often as important in the selection process as the mechanical properties of the alloy. A common solution to a corrosion problem is to substitute and alloy with greater corrosion resistance for the alloy that has corroded.
Coatings Coatings for corrosion protection can be divided into two broad groups—metallic and nonmetallic (organic and inorganic). With either type of coating the intent is the same, that is, to isolate the underlying metal from the corrosive media. Metallic Coatings. The concept of applying a more noble metal coating on an active metal takes advantage of the greater corrosive resistance of the noble metal. An example of this application is tin-plated steel. Alternatively, a more active metal can be applied, and in this case the coating corrodes preferentially, or sacrificially, to the substrate. An example of this system is galvanized steel, where the sacrificial zinc coating corrodes preferentially and protects the steel. Organic Coatings. The primary function of organic coatings in corrosion protection is to isolate the metal from the corrosive environment. In addition to forming a barrier layer to stifle corrosion, the organic coating can contain corrosion inhibitors. Many organic coating formulations exist, as do a variety of application processes to choose from for a given product or service condition.
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Inorganic coatings include porcelain enamels, chemical-setting silicate cement linings, glass coatings and linings, and other corrosionresistant ceramics. Like organic coatings, inorganic coatings for corrosion applications serve as barrier coatings. Some ceramic coatings, such as carbides and silicides, are used for wear-resistant and heatresistant applications, respectively.
Inhibitors Just as some chemical species (e.g., salt) promote corrosion, other chemical species inhibit corrosion. Chromates, silicates, and organic amines are common inhibitors. The mechanisms of inhibition can be quite complex. In the case of the organic amines, the inhibitor is adsorbed on anodic and cathodic sites and stifles the corrosion current. Other inhibitors specifically affect either the anodic or cathodic process. Still others promote the formation of protective films on the metal surface. The use of inhibitors is favored in closed systems where the necessary concentration of inhibitor is more readily maintained. The increased use of cooling towers stimulated the development of new inhibitor/ water-treatment packages to control corrosion and biofouling. Inhibitors can be incorporated in a protective coating or in a primer for the coating. At a defect in the coating, the inhibitor leaches from the coating and controls the corrosion.
Cathodic Protection Cathodic protection suppresses the corrosion current that causes damage in a corrosion cell and forces the current to flow to the metal structure to be protected. Thus, the corrosion or metal dissolution is prevented. In practice, cathodic protection can be achieved by two application methods, which differ based on the source of the protective current. An impressed-current system uses a power source to force current from inert anodes to the structure to be protected. A sacrificial-anode system uses active metal anodes, for example, zinc or magnesium, which are connected to the structure to provide the cathodic-protection current.
Design The application of rational design principles can eliminate many corrosion problems and greatly reduce the time and cost associated with corrosion maintenance and repair. Corrosion often occurs in dead spaces or crevices where the corrosive medium becomes more corrosive. These areas can be eliminated or minimized in the design process. Where stress-corrosion cracking is possible, the components can be designed to operate at stress levels below the threshold stress for cracking.
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9
Where corrosion damage is anticipated, design can provide for maximum interchangeability of critical components and standardization of components. Interchangeability and part standardization reduce the inventory of parts required. Maintenance and repair can be anticipated, and easy access can be provided. Furthermore, for the large items that are critical to the entire operation, such as primary pumps or large fans, redundant equipment is installed to permit maintenance on one unit while the other is operating. These practices are a sampling of rational design principles.
Opportunities in Corrosion Control The massive costs of corrosion provide many opportunities to users, manufacturers, and suppliers. Opportunities exist to reduce corrosion costs and the risks of failure, and to develop new, expanded markets. Examples of these opportunities and the means to implement a program to capitalize on the opportunities are presented in Table 1. The costs of corrosion vary considerably from industry to industry; however, substantial savings are achievable in most industries. The first step in any cost-reduction program is to identify and quantify the present costs of corrosion. Based on this analysis and a review of the present status of corrosion control in the industry, priorities can be determined and the most rewarding cost-reduction projects pursued. Risk of corrosion failure can be lowered in the producer’s facility and in its products. Both process and products can be analyzed to identify the areas where corrosion failures can occur. Once identified, the risk of failure can be evaluated from the perspectives of impact on safety, product liability, avoidance of regulation, and loss of goodwill. Where risks Table 1
Opportunities in corrosion control
Opportunity
Reduce corrosion costs
Lower risk of failure
Develop new and expanded markets
Examples
Lower maintenance and repair costs Extended useful lives of equipment and buildings Reduction of product loss from corrosion damage Safety Product liability Avoidance of regulation Loss of goodwill
Coatings Alloys Inhibitors Corrosion monitors
Implementation
Identify all corosion costs by review of total processes, equipment, and buildings Quantify corrosion costs Implement plan to reduce costs Review process and products for exposure to risk Evaluate risk and consequences of failure Lower exposure by technology change Apply emerging technology Develop competitive advantage by more corrosion-resistant product Transfer existing technology to other industries
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are too great, technological changes can be implemented to reduce the risk. Evaluation also can identify areas where technological advances are required in the industry. Increased consumer awareness of corrosion provides a competitive advantage for products with improved corrosion resistance. Through the application of existing or emerging technologies to products or services, advances are being made in all methods for corrosion control: material selection, coatings, inhibitors, cathodic protection, and design. Market opportunities are to be found in the transfer of existing technology to other industries.
The Economic Impact of Corrosion Corrosion of metals costs the U.S. economy almost $300 billion per year at current prices. Approximately one-third of these costs could be reduced by broader application of corrosion-resistant materials and the application of best corrosion-related technical practices. These estimates result from a recent update of findings of the 1978 study Economic Effects of Metallic Corrosion in the United States. The study was performed by Battelle Columbus Laboratories and the National Institute of Standards and Technology (NIST) and published in April 1995. The original work, based upon an elaborate model of more than 130 economic sectors, found that in 1975, metallic corrosion cost the United States $82 billion, or 4.9% of its gross national product (GNP). It was also found that 60% of that cost was unavoidable. The remaining $33 billion (40%) was incurred by failure to use the best practices then known. These were called “avoidable” costs. Over the last two decades, economic growth and price inflation have increased the GNP more than fourfold. If nothing else had changed, the costs of metallic corrosion would have risen to almost $350 billion annually by 1995, $139 billion of which would have been avoidable. However, 20 years of scientific research and technological change, much of which was initiated because of the 1978 study, have affected these costs. The Battelle panel updated the earlier results by judgmentally evaluating two decades of corrosion-related changes in scientific knowledge and industrial practices. In the original study, almost 40% of the 1975 metallic corrosion costs were incurred in the production, use, and maintenance of motor vehicles. No other sector accounted for as much as 4% of the total, and most sectors contributed less than 1%. The aircraft sector, for instance, was one of the next largest contributors and accounted for just more than 3%. Pipelines, a sector to which corrosion is a recognized problem, accounted for less than 1% of the total cost.
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The panel found that the automotive sector probably had made the greatest anticorrosion effort of any single industry. Advances have been made in the use of stainless steels, coated metals, and more protective finishes. Moreover, several substitutions of materials made primarily for reasons of weight reduction have also reduced corrosion. Also, the panel estimates that 15% of previously unavoidable corrosion costs can be reclassified as avoidable. The industry is estimated to have eliminated some 35% of avoidable corrosion by improved practices. In examining the aircraft, pipeline, and shipbuilding sectors, the panel reported that both gains and losses have occurred, most of them tending to offset each other. For instance, in many cases, the use of more expensive materials has reduced the need for corrosion-related repairs or repainting. Overall, it was thought that for the U.S. economy other than in motor vehicle and aircraft applications, total corrosion costs have been reduced by no more than 5% with a further reduction of unavoidable costs by about 2%. The updated study shows that the total 1995 cost of metallic corrosion was reduced (from what it would have been in 1975 terms) by some 14%, or to 4.2% of the GNP. Avoidable corrosion, which was 40% of the total, is now estimated to be 35% but still accounts for slightly more than $100 billion per year. This figure represents the annual cost to the economy, which can be reduced by broader application of corrosionresistant materials, improvement in corrosion-prevention practices, and investment in corrosion-related research. Table 2 compares the results of the 1978 and 1995 Battelle/NIST studies. Factors Influencing Corrosion. Some of the factors that influence corrosion and its costs are shown in Fig. 4. Corrosion costs are reduced by the application of available corrosion technology, which is supTable 2
Cost of metallic corrosion in the United States Billions of U.S. dollars
Industry
1975
1995
82.0 33.0
296.0 104.0
31.4 23.1
94.0 65.0
3.0 0.6
13.0 3.0
47.6 9.3
189.0 36.0
All industries Total Avoidable Motor vehicles Total Avoidable Aircraft Total Avoidable Other industries Total Avoidable
Source: Economic Effects of Metallic Corrosion in the United States, Battelle Columbus Laboratories and the National Institute of Standards and Technology (NIST), 1978, and Battelle estimates
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Corrosion: Understanding the Basics
ported by technology transfer. New and improved corrosion technology results from research and development. The proper application of methods to control corrosion (e.g., coatings, inhibitors, and cathodic protection) reduces the cost of corrosion. The costs of corrosion tend to increase with such factors as deferred maintenance and extended useful lives of buildings and equipment. Increased corrosion costs are often realized when higher-performance specifications and more hostile environments are encountered. Finally, increased corrosion costs result from government regulations that prohibit the use of time-honored methods of protection because of safety or environmental damage. For example, in an effort to reduce smog, the elimination of lead-based paints on houses and bridges, chromate inhibiting paints on aircraft, and oil-based paints throughout industry has had severe repercussions. Substitute water-based paints have not, in many cases, afforded equivalent corrosion protection. Cost Elements. Although costs vary in relative significance from industry to industry, several generalized elements combine to make up the total cost of corrosion. Some are readily recognized; others are less recognizable. In manufacturing, corrosion costs are incurred in the product development cycle in several ways, beginning with the materials, energy, labor, and technical expertise required to produce a product. For example, a product can require painting for corrosion protection. A corrosionresistant metal can be chosen in place of plain carbon steel, and technical services can be required to design and install cathodic protection on a product. Additional heat treatment can be needed to relieve stresses for protection against stress-corrosion cracking. Other operating costs are affected by corrosion as well. Corrosion inhibitors, for example, often must be added to water treatment systems. Applied current technology More hostile environments
Deferred maintenance
Increased performance requirements
Technology transfer
Fig. 4
Corrosion costs
Extensions of useful life
Environmental regulations
Research and development
Factors which increase or decrease the costs of corrosion
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Portions of maintenance and repair costs can be attributed to corrosion, and corrosion specialists are often employed to implement corrosioncontrol programs. Capital costs also are incurred because of corrosion. The useful life of manufacturing equipment is decreased by corrosion. For an operation that is expected to run continuously, excess capacity is required to allow for scheduled downtime and corrosion-related maintenance. In other instances, redundant equipment is installed to enable maintenance on one unit while processing continues with another unit. For the end user or consumer, corrosion costs are incurred for purchases of corrosion prevention and control products, maintenance and repair, and premature replacement. The original Battelle/NIST study identified ten elements of the cost of corrosion: · · · · · · · · · ·
Replacement of equipment or buildings Loss of product Maintenance and repair Excess capacity Redundant equipment Corrosion control Technical support Design Insurance Parts and equipment inventory
Table 3 lists examples under each of these categories. Replacement, loss of product, and maintenance and repair are fairly straightforward. Excess capacity is a corrosion cost if downtime for a plant scheduled for continuous operation could be reduced were corrosion not a factor. This element accounts for extra plant capacity (capital stock) maintained because of corrosion. Redundant equipment accounts for additional plant equipment (capital stock) required because of corrosion. Specific critical components such as large fans and pumps are backed up by identical items to allow processing to continue during maintenance for corrosion control. The costs of corrosion control are straightforward, as are the technical support (engineering, research and development, and testing) costs associated with corrosion. Corrosion costs associated with design are not always as obvious. The last two cost elements, insurance and inventory, can be significant in specific cases. In addition to these ten categories, other less quantifiable cost factors, such as loss of life or loss of goodwill because of corrosion, can have a major impact. Single, catastrophic failures—for example, a corrosion-
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Corrosion: Understanding the Basics
Table 3
Elements of cost of corrosion
Element of cost
Replacement of equipment or buildings Loss of product
Maintenance and repair
Redundant equipment Corrosion control Inhibitors Organic coatings
Metallic coatings Cathodic protection Technical support
Design Material of construction for structural integrity Material of construction Corrosion allowance Special processing for corrosion resistance Insurance
Parts and equipment inventory
Example
Corroded pressure vessel Corrosion leak Corrosion contamination of product Corrosion during storage Repair corroded corrugated metal roof Weld overlay of chemical reaction tank Repair pump handling corrosive slurry—erosion and corrosion Scheduled downtime for plant in continuous operation, for example, petroleum refinery Installation of three large fans where two are required during operation Injection of oil wells Coal tar on exterior of underground pipeline Paint on wooden furniture Topcoat on automobile—aesthetics and corrosion Zinc-rich paint on automobile Galvanized steel siding Chrome-plated faucets—aesthetics and corrosion Cathodic protection of underground pipelines Corrosion-resistant alloy development Materials selection Corrosion monitoring and control Stainless steel for corrosive applications Stainless steel for high-temperature mechanical properties High alloy to prevent corrosion products contamination, for example, drug industry Thicker wall for corrosion Stress relief, shot peening, special heat treatment (e.g., Al alloys) for corrosion Portion of premiums on policy to protect against loss because of corrosion (to cover charge of writing and administering policy, not protection amount) Pumps kept on hand for maintenance, for example, chemical plant inventory
Source: Ref 1
induced leak in an oil pipeline, with resulting loss of product and environmental contamination—can result in costly damage that is difficult to either assess or repair as well as massive legal penalties as “punative damage.”
Sources of Information Sources of information pertaining to corrosion and corrosion prevention are quite varied and include the following: · · · ·
Texts, reference books, and journals Videos and home study courses Software products Computerized databases
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· Metals producers · Trade associations and technical societies · Consultants
Titles of several widely used textbooks on corrosion and a comprehensive bibliography relevant to corrosion are provided at the conclusion of this chapter (see the Selected References). Complementing print products are video training courses that are available from ASM International (formerly the American Society for Metals) and NACE International (formerly the National Association of Corrosion Engineers). Reference works that list corrodents in alphabetical order and give information for a variety of metallic and nonmetallic materials are particularly useful. Some provide only qualitative information such as “Resistant,” “Unsatisfactory,” etc., but others can give a more specific indication of the general corrosion rate. An example of the latter approach is Corrosion Resistance Tables: Metals, Nonmetals, Coatings, Mortars, Plastics, Elastomers and Linings, and Fabrics published by Marcel Dekker. In the Corrosion Data Survey—Metals and its companion volume, Corrosion Data Survey—Nonmetals, published by NACE International, the corrosion rate of a given material is plotted against temperature and corrodent concentration. Electronic versions of these products are also described in Chapter 8. A number of technical journals on the subject of corrosion exist. Examples include Corrosion, and Materials Performance, published by NACE International, and Oxidation of Metals, published by Plenum Publishing Corp. Journals covering corrosion science and technology can also be found in numerous other metallurgical, surface engineering (coating), chemical, and electrochemical publications. The Source Journals in Metals & Materials, available in print or electronic format from Cambridge Scientific Abstracts (Beachwood, OH) lists dozens of journals devoted to corrosion. Producers of metals and alloys publish considerable product data and educational information, as do trade associations such as the Nickel Development Institute, the Aluminum Association, the Copper Development Association, and the Specialty Steel Industry of North America. Addresses for these and other associations and societies are listed in the appendix to this chapter. Research organizations such as the LaQue Center for Corrosion Technology (Wrightsville Beach, NC) and the Electric Power Research Institute (Palo Alto, CA) also provide extensive corrosion information. Several technical societies are involved with corrosion work. They serve as a source of technical literature, standards, reports, and software. They also sponsor technical symposia and have technical committees that cover a broad spectrum of corrosion problems. In the United States, the primary society devoted to corrosion is NACE Inter-
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Table 4 Committee
T-1 T-2 T-3 T-5 T-6 T-7 T-8 T-9 T-10 T-11 T-14
Table 5 Subcommittee
G01.02 G01.03 G01.04 G01.05 G01.06 G01.07 G01.08 G01.09 G01.10 G01.11 G01.12 G01.14 G01.99.01
NACE International technical committees Activity
Corrosion control in petroleum production Energy technology Corrosion science and technology Corrosion problems in the process industries Protective coatings and linings Corrosion by waters Refining industry corrosion Military, aerospace, and electronics equipment corrosion control Underground corrosion control Corrosion and deterioration of the infrastructure Corrosion in the transportation industry
ASTM committee G-1 on corrosion of metals Activity
Terminology Computers in corrosion Atmospheric corrosion Laboratory corrosion tests Stress-corrosion cracking and corrosion fatigue Galvanic corrosion Corrosion of nuclear materials Corrosion in natural waters Corrosion in soils Electrochemical measurements in corrosion testing In-plant corrosion tests Corrosion of reinforcing steel Corrosion of implant materials
national. NACE was formed in 1943 with the aim of assisting the public and industry in the use of corrosion prevention and control to reduce the billions of dollars lost each year caused by corrosion. Table 4 lists NACE technical committees. NACE also sponsors a yearly international congress on corrosion. ASTM (formerly the American Society for Testing and Materials) is also very active in the field of corrosion. The main committee is G-1 on corrosion of metals. Its scope is “the promotion of knowledge, the stimulation of research, the collection of engineering data, and the development of standard test methods, practices, guides, classifications, specifications and terminology relating to corrosion and methods for corrosion-protection of metals.” A list of the subcommittees in G-1 is shown in Table 5. Other societies having interests in corrosion are the American Institute of Mining, Metallurgical, and Petroleum Engineers; the American Petroleum Institute; the Electrochemical Society; the American Institute of Chemical Engineers; the American Welding Society; ASM International; the American Society of Mechanical Engineers; the Society for Protective Coatings (formerly the Steel Structures Painting Council); and SAE International (formerly the Society of Automotive Engineers). Most of these societies have symposia on corrosion at their various meetings.
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Appendix: Addresses of Trade Associations and Technical Societies Involved with Corrosion
Aluminum Association, Inc. 900 19th St., NW Suite 300 Washington, DC 20006 American Institute of Mining, Metallurgical, and Petroleum Engineers (AIME) 345 E. 47th St., 14th Floor New York, NY 10017 American Iron and Steel Institute (AISI) 1101 17th St., NW Suite 1300 Washington, DC 20036-4700 American National Standards Institute (ANSI) 11 W. 42nd St., 13th Floor New York, NY 10036 American Petroleum Institute (API) 1220 L St., NW Washington, DC 20005 American Society of Mechanical Engineers (ASME) 345 E. 47th St. New York, NY 10017 American Welding Society (AWS) 550 N.W. LeJeune Rd. Miami, FL 33126 ASM International 9639 Kinsman Rd. Materials Park, OH 44073-0002
ASTM 100 Barr Harbor Dr. W. Conshohocken, PA 19428-2959 Canadian Institute of Mining, Metallurgy, and Petroleum (CIM) Xerox Tower Suite 2110 3400 de Maisonneuve Blvd., W. Montreal, QC Canada, H3Z 3B8 Canadian Standards Association (CSA) 178 Rexdale Blvd. Rexdale, ON Canada M9W 1R3 Copper Development Association (CDA) 260 Madison Ave. New York, NY 10016 International Cadmium Association 12110 Sunset Hills Rd. Suite 110 Reston, VA 22090 International Copper Association Ltd. 260 Madison Ave. New York, NY 10016 International Lead Zinc Research Organization, Inc. (ILZRO) 2525 Meridian Parkway P.O. Box 12036 Research Triangle Park, NC 27709
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International Magnesium Association (IMA) 1303 Vincent Place Suite 1 McLean, VA 22101 International Titanium Association (ITA) 1781 Folsom St. Suite 100 Boulder, CO 80302-5714 Lead Industries Association, Inc. 295 Madison Ave. New York, NY 10017 Materials Technology Institute of the ChemicalProcessIndustries,Inc.(MTI) 1570 Fishinger Rd. Columbus, OH 43221 NACE International P.O. Box 218340 Houston, TX 77218-8340
SAE International 400 Commonwealth Dr. Warrendale, PA 15096-0001 Society for the Advancement of Materials and Processing Engineering (SAMPE) P.O. Box 2459 Covina, CA 91722 Specialty Steel Industry of North America (SSINA) 3050 K St., NW Suite 400 Washington, DC 20007 Steel Founders’ Society of America (SFSA) Cast Metals Federation Building 455 State St. Des Plaines, IL 60016
National Institute of Standards and Technology (NIST) Gaithersburg, MD 20899
The Society for Protective Coatings (SSPC) 40 24th St. 6th Floor Pittsburgh, PA 15222-4643
Nickel Development Institute (NiDI) 214 King St., W. Suite 510 Toronto, ON Canada M5H 3S6
The Metallurgical Society (TMS-AIME) 420 Commonwealth Dr. Warrendale, PA 15086-7514
References 1. J.H. Payer et al., Mater. Perform., Vol 19 (No. 9), June 1980, p 19–20
Selected References · A Glossary of Corrosion-Related Terms Used in Science and Industry, M.S. Vukasovich, SAE International, 1995
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· S.A. Bradford, Practical Self-Study Guide to Corrosion Control, Casti Publishing, 1998 · Corrosion, Vol 13, ASM Handbook, ASM International, 1987 · Corrosion and Corrosion Protection Handbook, 2nd ed., P.A. Schweitzer, Ed., Marcel Dekker, 1989 · Corrosion Basics—An Introduction, L.S. Van Delinder, Ed., NACE International, 1984 · Corrosion Data Survey—Metals Section, 6th ed., D.L. Graver, Ed., NACE International, 1985 · Corrosion Data Survey—Nonmetals Section, 5th ed., NACE International, 1975 · Corrosion Engineering Handbook, P.A. Schweitzer, Ed., Marcel Dekker, 1996 · Corrosion Resistance Tables, 4th ed., 3-volume set, P.A. Schweitzer, Ed., Marcel Dekker, 1995 · Corrosion-Resistant Materials Handbook, 4th ed., D.J. DeRenzo, Ed., Noyes, 1985 · Corrosion Source Book, S.K. Coburn, Ed., American Society for Metals, 1984 · R.W. Drisko and J.F. Jenkins, Corrosion and Coatings: An Introduction to Corrosion for Coatings Personnel, The Society for Protective Coatings, 1998 · E.D. During, Corrosion Atlas, 3rd ed., Elsevier Scientific Publishers, 1997 · M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill Book Company, 1986 · Handbook of Corrosion Data, 2nd ed., B. Craig and D. Anderson, Ed., ASM International, 1995 · D.A. Jones, Corrosion Principles and Prevention of Corrosion, 2nd ed., Prentice Hall, 1996 · P. Marcus and J. Oudar, Corrosion Mechanisms in Theory and Practice, Marcel Dekker, 1995 · E. Mattson, Basic Corrosion Technology for Scientists and Engineers, 2nd ed., The Institute of Materials, 1996 · NACE Corrosion Engineer’s Reference Book, 2nd ed., R.S. Treseder, R. Baboian, and C.G. Munger, Ed., NACE, 1991 · P.A. Schweitzer, Encyclopedia of Corrosion Technology, Marcel Dekker, 1998 · P.A. Schweitzer, What Every Engineer Should Know About Corrosion, Marcel Dekker, 1987 · J.C. Scully, The Fundamentals of Corrosion, 3rd ed., Pergamon Press, 1990
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· D. Talbot and J. Talbot, Corrosion Science and Technology, CRC Press, 1997 · H.H. Uhlig and R.W. Revie, Corrosion and Corrosion Control, 3rd ed., John Wiley & Sons, 1985
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Corrosion: Understanding the Basics J.R. Davis, editor, p21-48 DOI: 10.1361/cutb2000p021
CHAPTER
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2
Basic Concepts Important to Corrosion BASIC CONCEPTS important to understanding corrosion that are addressed in this chapter include the following: · · · · · ·
Three possible behaviors of a metal when immersed in a solution Four requirements of a corrosion cell Important metallurgical factors that influence corrosion behavior Inherent tendency of a metal to corrode, that is, reactivity Tendency of metals to form corrosion products Important solution characteristics with respect to corrosion, including conductivity, acidity/alkalinity, oxidizing power, and solubility · Determination of corrosion rates and corrosion rate allowances These principal concepts are referred to throughout this book to assist the reader in understanding corrosion phenomena and methods of controlling corrosion. Information pertaining to important electrochemical and thermodynamic reactions is in Chapter 3.
Behavior of a Metal in an Environment When a metal is immersed in an environment, the metal can behave in one of three ways. These behaviors are shown schematically in Fig. 1, which represents a metal partially immersed in a corrosive environment.
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Corrosion: Understanding the Basics
Immune Behavior. One possibility shown in Fig. 1 is that the metal is immune in an environment. Metals known to display this immunity are called noble metals and include, for example, gold, silver, and platinum. For a combination of metal and environment resulting in immune behavior, there is no reaction of the metal, and there is no corrosion of the metal. If the metal is weighed prior to immersion in the solution and then reweighed after the exposure period, there is no weight loss of the metal. Immune behavior results from the metal being thermodynamically stable in the particular environment; that is, the corrosion reaction does not occur spontaneously. Active Behavior. Another possible behavior is that the metal corrodes. A metal’s behavior is described as active when it corrodes in the solution. When active behavior is observed, the metal dissolves in solution and forms soluble, nonprotective corrosion products. Corrosion or dissolution of the metal continues in this solution because the corrosion products do not prevent subsequent corrosion. Active corrosion is characterized by high weight loss of the metal. If the metal sample is weighed prior to immersion in the solution and then reweighed after the exposure period, a significant weight loss is measured. Passive Behavior. With the third behavior, the metal corrodes but a state of passive behavior is observed. On immersion of the metal in the solution there is a reaction, and the metal does corrode; however, an insoluble, protective corrosion-product film is formed. This thin (~30Å) protective film, also referred to as a passive film, slows the reaction rate to very low levels. The corrosion resistance when dealing with passive behavior depends on the integrity of the protective film. If the passive film is broken or dissolves, then the metal can revert to active behavior and rapid dissolution can occur. For example, an iron sample immersed in either concentrated (70%) or dilute (water added) nitric acid exhibits no reaction (immune behavior) or passive behavior. However, if the iron is scratched with a glass rod or if the beaker holding the sample/ solution is shaken violently so that the sample strikes the sides, a violent reaction occurs. The iron quickly goes into solution and large volumes of nitrogen-bearing gases are released. Some examples of metals that exhibit passivity are iron, chromium, titanium, nickel, and alloys containing these metals (most notably stain-
Fig. 1
Three behaviors of metal in an environment
Basic Concepts Important to Corrosion
23
less steels). Passivation is generally associated with oxidizing media. (Discussion on the oxidizing power of a solution follows.) It is important to realize that a passive film is unlike a coat of paint— though for many practical purposes it can appear to behave as such. This point is of practical significance with respect to the so-called “passivation treatments.” These treatments are commonly used for stainless steels. They involve immersing the steel in an oxidizing solution such as nitric acid for approximately half an hour. The prime purpose of such a treatment is not (as is commonly believed) to form a passivation film, but rather to clean the steel—to remove surface inclusions, iron particles, etc., that might act as nucleation sites for attack in future service. In other words, the passivation treatment is only useful in that it creates surface conditions (cleanliness) that can make the stainless steel more amenable to maintaining its passivation. What is the Desired Corrosive Behavior? From a corrosioncontrol standpoint, the desired behavior is either immune or passive, while the behavior to be avoided is active. Immune behavior is the most desirable, because corrosion protection does not depend on the stability of protective films. Most engineering alloys, however, are passive in their applications and thus depend on the integrity of the passive film. Where the environment becomes more corrosive, passive metals tend to exhibit localized forms of corrosion, that is, pitting, stress-corrosion cracking, and crevice corrosion. This results because the bulk of the alloy surface remains protected by the passive film, but rapid corrosion occurs in those areas where the film has broken down. Only the most noble metals exhibit immune behavior in a wide variety of corrosive environments. In most cases, it is not practical to use these materials for engineering applications because of their high costs and strength limitations. While the behavior exhibited by metals is dependent on the corrosive environment to which the metals are exposed, some characteristic behaviors are exhibited. As mentioned above, gold, platinum, and silver typically exhibit noble or immune behavior. Sodium, potassium, and magnesium are active in nearly all aqueous environments. Titanium and tantalum are passive in a wide range of aqueous environments. Aluminum and zinc are very reactive metals and often exhibit active behavior; however, in some important environments they form stable, passive films. The important characteristics of metals and solutions that determine the type of behavior observed are discussed in the following sections of this chapter.
The Four Requirements of a Corrosion Cell There are four requirements for an electrochemical corrosion cell. These are shown schematically in Fig. 2, where an anode and a cathode
24
Corrosion: Understanding the Basics
on the metal surface in contact with the solution are indicated. The anode and cathode are connected through the solution by an ionic current path, and they are connected through the metal by an electronic path. An electrochemical reaction involves the transfer of electrons from one species to another, causing direct current flow through the corrosion cell. The anode is generally where the corrosion occurs. This is the location on the metal surface where metal atoms go into solution as metal ions and weight loss occurs. The direct current going through the corrosion cell enters the solution at the anode. The reactions at the anode are referred to as anodic and are oxidation reactions; that is, electrons are generated. At the cathode, no corrosion occurs and no weight loss occurs. There are, however, reactions occurring that are just as important to the operation of the corrosion cell as the anodic reactions. These reactions are cathodic or reduction reactions; that is, electrons are consumed. The direct current flowing through the corrosion cell enters that metal at the cathode. The direct current of the corrosion cell moves through the solution by an ionic path. Current flows from the anode to the cathode by the movement of charged ions in the solution. Positively charged ions, or cations, move from the anode to the cathode, and negatively charged ions, or anions, move from the cathode toward the anode. This movement of charged ions in the solution is the vehicle for current flow through this portion of the corrosion cell. The direct current moves through the metal of the corrosion cell by an electronic path. Electrons generated at the anode by oxidation reactions move to the cathode, where they are consumed by reduction reactions. Current is, by convention, the flow of positively charged particles. Thus, the current (positive charges) flows conceptually from the cathode to the anode. These four requirements make up the corrosion cell. Metal atoms going into solution at the anode result in corrosion and the generation of electrons at the anode. Current flows from the anode to the cathode by the movement of charged particles. At the cathode, reactions occur that consume electrons, that is, reduction reactions, and the electrons generIonic current path
Anode
Cathode
Electronic path
Fig. 2
Four requirements of an electrochemical corrosion cell
Basic Concepts Important to Corrosion
25
ated at the anode are consumed. The electronic path is the path by which electrons move from the anode to the cathode. The corrosion rate is controlled by the net balance among all of these components of the corrosion cell. The dissolution (oxidation) at the anode can only proceed as quickly as the electrons generated there can be consumed by reduction reactions at the cathode. If the reduction reactions are slowed down, this in turn slows down the dissolution reactions. Resistance in the ionic current path or the electronic current path will slow down the corrosion reaction by limiting the amount of current that can flow through the corrosion cell. Elimination of any of the four requirements for the corrosion cell stops the corrosion reaction. If the anodes are removed or made inactive, no metal dissolution can occur. An effective control of corrosion is realized by the elimination of the cathodes. If there is no place for the consumption of electrons generated by the corrosion reaction, there is no corrosion reaction. Elimination of the ionic current path also stops corrosion. There is no means for the transfer of electrical charge from the anodes to the cathodes. A practical example of this form of corrosion control is the removal of the electrolyte in a corrosion cell. This can be done by completely drying the metal surface. If there is no moisture on the surface for the formation of an ionic current path, there is no aqueous corrosion. Similarly, elimination of the electronic path between the anode and cathode also eliminates corrosion. If we are dealing with galvanic corrosion, a corrosion reaction driven by two dissimilar metals, the galvanic corrosion can be eliminated by electrically isolating the two metals. There then is no path by which electrons can be transferred from the anode to the cathode, and the two metals do not affect each other.
Metal Characteristics Important to Corrosion A knowledge of what metals are and how they behave is essential to the understanding of corrosion. In this section, the important characteristics of metals with respect to corrosion are identified. For metals, the metallurgical characteristics, inherent reactivity, and tendency to form insoluble corrosion products all greatly affect their corrosion behavior.
Metallurgical Characteristics Crystal Structure. Metals are crystallographic in nature; that is, the metal atoms are arranged in an ordered and structured manner throughout the metal crystal. This can be demonstrated by considering each metal atom in the crystal as a sphere. Based on this, models can be built to represent various metal crystal structures. Three such metal crystal
26
Corrosion: Understanding the Basics
structures are shown in Fig. 3. The atomic packing of atoms in metal crystals with face-centered cubic (fcc), hexagonal close-packed (hcp), and body-centered cubic (bcc) structures are shown. In each structure, the metal atoms have a very well-defined, repeatable, and orderly relationship to one another. The metal crystal can be assembled by putting together layers of planes to build up the overall volume of the crystal. Each plane has the identical arrangement of metal atoms within it. The surface appearance of a metal crystal on the atomic scale depends on the angle of the planes intercepting that surface. This is shown schematically in Fig. 4. If the surface is directly along the angle of the planes (alpha = 0), then only a single atomic plane is exposed along the surface. As the angle of interception of planes with the surface increases, more and more edges of the planes are exposed. The angles at which the planes intercept the surface affect the reactivity of the metal and its resistance to corrosion because the binding energy of the “end” atoms is less than that of atoms in the plane.
(a)
(b)
(c)
Fig. 3
Unit cells and atom positions for (a) face-centered cubic, (b) hexagonal close-packed, and (c) body-centered cubic unit cells. The positions of the atoms are shown as dots at the left of each pair of drawings, while the atoms themselves are shown close to their true effective size by spheres or portions of spheres at the right of each pair.
Basic Concepts Important to Corrosion
27
Grain Boundaries. Most materials used in service are not single crystals but are in fact made up of many individual crystals or grains. In each grain the metal has planes characteristic of the crystal structure of that metal. The planes from grain to grain are not in the same orientation. This gives rise to grain boundaries between the adjoining crystals. Grain boundaries are shown schematically in Fig. 5. Essentially, the grain boundaries are the area of transition from orientation within one grain to orientation in the neighboring grain. A micrograph of the grain boundaries in a low-carbon steel is shown in Fig. 6. Grain boundaries are sites of structural discontinuity, and they can also have microstructural and chemical differences with respect to the bulk grains. These discontinuities and differences can affect the corrosion behavior of the metal. Alloying and Multiphase Structures. While pure metals have many applications, mixtures of several different elements are worked with much more commonly. These intentional mixtures of elements to obtain desirable properties are called alloys. Two or more elements mixed together give rise to metals with a wide variety of properties not
Fig. 4
Crystallographic planes intersecting the surface at different angles
Crystalline grains, zones of near-perfect fit
Grain boundary, zone of misfit
Microstructure
Fig. 5
Atomic arrangement
Schematic diagram of grain boundaries in a metal
28
Corrosion: Understanding the Basics
available using single elements. The microstructure resulting from the mixture of two elements can vary widely. In some cases the two elements are completely soluble and a homogeneous, single-phase structure is exhibited. Other metals have only limited solubility, and mixtures of these elements result in multiphase materials. One of the most useful tools for studying the effects of alloying on microstructure is the phase diagram. This is a graph that plots the phase stability relation between various compositions of one metal in another as a function of temperature. In other words, it shows all possible phases of the various possible alloy mixtures and the temperatures at which these phases exist. An excellent introduction to the use and understanding of alloy phase diagrams can be found in Alloy Phase Diagrams, Volume 3 of the ASM Handbook (see pages 1·1 to 1·29). An example of a phase diagram showing complete solid solubility for the copper-nickel system is shown in Fig. 7. A copper-nickel alloy will be a homogeneous, single phase at any percentage of nickel from pure copper across the diagram to pure nickel. The alloy has grain boundaries in the regions between single crystals of different orientation; however, there is only a single phase of constant composition in all grains. The copper-silver phase diagram with limited solid solubility is shown in Fig. 8. Copper is very sparingly soluble in silver, and silver is very sparingly soluble in copper. This results in two-phase structures being exhibited across nearly the entire range of composition of copper and silver alloys. There is a phase consisting of nearly pure copper and another phase consisting of nearly pure silver. The ratio of amounts of the two phases varies depending on the relative amounts of copper and silver. The copper-silicon phase diagram shown in Fig. 9 exhibits many different phases with increasing silicon content from pure copper (0% Si)
Fig. 6
Ferrite grains and grain boundaries in a low-carbon ferritic sheet steel etched with 2% nital. 300×
Basic Concepts Important to Corrosion
29
to 14 wt% Si. The phases that are present and their relative amounts depend on the composition of the copper-silicon alloy and also on the heat treatment of the alloy. There are literally thousands of examples of micrographs showing the distribution and morphology of multiple-phase microstructures in the ASM Handbook series as well as in other ASM publications. Two notable books are Metallography and Microstructures, Vol 9, ASM Handbook, which deals with metallographic preparation and microstructural interpretation of industrial alloys, and the Metals Handbook Desk Edition, Second Edition. Of particular note in the latter publication is the article “Structure/Property Relationships in Irons and Steels”
Liquid (L) (L + )
Fig. 7
Copper-nickel phase diagram with complete solid solubility. The diagram consists of two single-phase fields separated by a two- phase field (L + a). The boundary between the liquid field (L) and the two- phase field is called the liquidus; that between the two-phase field and the solid field (a) is the solidus.
Fig. 8
Copper-silver phase diagram with limited solubility
30
Corrosion: Understanding the Basics
(pages 153 to 173), which shows many of the various structures possible in iron-base alloys. Relationship between Microstructure and Corrosion. The important consideration within the context of this corrosion course is that many alloys are not homogeneous, pure materials, but rather are a mixture of multiple phases. Each phase has its characteristic crystallographic structure and chemical composition. When these structures are then exposed to a corrosive environment, it is not surprising that the different phases exhibit different corrosion behaviors. This leads to preferential corrosion of specific constituents of the alloy. This relationship between microstructure and corrosion behavior is demonstrated in Fig. 10 and 11. Figure 10 shows the microstructure of three different alloys. Alloy 1 is nearly pure A, alloy 2 is A with modest amounts of B, and alloy 3 is a B-rich alloy of A and B. The microstructure of alloy 1 is a single phase of alpha (a), with complete solubility of B within the alpha phase. Alloy 2 has B-additions beyond the solubility limit, and a two-phase structure results. The microstructure is made up of small islands of beta (b) phase distributed throughout a continuous matrix of alpha phase. The microstructure of alloy 3 is a mixture of alpha phase and beta phase. Figure 11 shows the relationship of the corrosion behavior to the microstructures of the three alloys. In the first scenario, alpha is the more active phase and beta is more noble; that is, the corrosion resistance of alpha is less than the corrosion resistance of beta. A cross section through the alloy surface after exposure to a corrosive environment is represented below the diagram of each microstructure. For the case where alpha is the more active material, a uniform corrosion of alpha is observed for alloy 1. For alloy 2, the beta phase is nearly unattacked and
Fig. 9
Copper-silicon phase diagram with multiple solid phases
Basic Concepts Important to Corrosion
Fig. 10
31
Relationship of microstructure to the phase diagram
the alpha phase on either side of the beta particle is attacked. For alloy 3, again the alpha phase is attacked and the exposed beta phase is left essentially unattacked. For the scenario where alpha is noble (more corrosion resistant) and beta is active, the resulting surface profile is shown at the bottom of Fig. 11. In this case the alpha phase is not significantly corroded in alloy 1.
α active β noble
α noble β active
Fig. 11
Relationship of corrosion behavior to microstructure
32
Corrosion: Understanding the Basics
In alloy 2, the alpha phase remains unattacked, and significant dissolution of the beta phase occurs. Similarly, in alloy 3, the exposed grains of beta phase are attacked while the exposed alpha phase is left unattacked. This difference in the dissolution behavior of various phases in a multiphase structure provides the basis for optical metallography. In metallographic examinations, the specimen surface is polished to a mirror finish and then exposed to chemical etchants. The chemical solution preferentially attacks particular constituents of the alloy, and thus the microstructure of the alloy is revealed. Effect of Inclusions and Precipitates. Alloys are intentional mixtures of elements to gain desired properties. The microstructure of an alloy can contain multiple phases, and the distribution and amount of the second phase are controlled to develop desired properties, for example, increased strength or toughness. Other second-phase particles can be undesirable. Examples of undesirable precipitates are oxides and sulfides, which precipitate in the metal from dissolved oxygen and sulfur in the metal-producing process. This results in a distribution of inclusions (small particles of oxide, sulfide, etc.) throughout the alloy. When these inclusions are exposed at the metal surface to a corrosive environment, they can affect corrosion behavior. The effects of inclusions at the metal surface are shown schematically in Fig. 12. The uppermost figure represents an inclusion exposed at the metal surface prior to corrosion, and the lower diagrams indicate the behavior under different conditions. If the inclusion is active, that is, less corrosion resistant than the matrix, then the inclusion dissolves, leaving a hole or pit in the metal surface. If only portions of the inclusion are active, then the exposed portions are attacked, leaving the other portions intact. If the inclusion is noble (more corrosion resistant than the matrix), then accel-
Fig. 12
Effect of an inclusion at the metal surface
Basic Concepts Important to Corrosion
erated attack of the matrix adjacent to the noble inclusion can be observed. In other cases where the inclusion is inert to attack, accelerated corrosion adjacent to the inclusion can still occur because of a crevice generated between the inclusion and the matrix. The presence of precipitates with minor alloying elements and impurities can lead to problems, because phases with widely different electrochemical properties are then present. This can result in local variations in corrosion resistance. Also, the addition of alloying elements to improve the resistance to general, or uniform, corrosion can cause increased susceptibility to localized corrosion processes, such as pitting or intergranular corrosion. The metallurgical factors that can influence localized corrosion of stainless steels are shown in Fig. 13. The precipitation of nitrides or carbides along the grain boundaries of the stainless steel can result in depletion of chromium in regions surrounding the particles, which in turn leads to accelerated corrosion of the chromiumdepleted regions. Inclusions represented by a manganese sulfide particle at the surface can result in the initiation of pitting on the surface. Other phases in the stainless steel represented by delta ferrite and alpha prime can cause chemical inhomogeneities and/or structural inhomogeneities, which lead to the initiation of localized corrosion. The effect of mechanical deformation is shown by the generation of active slip steps, which can weaken the protective film on the stainless steel and lead to localized corrosion.
Fig. 13
Schematic of the microstructural variables that can influence the corrosion behavior of stainless steels
33
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Corrosion: Understanding the Basics
Effect of Conductivity. Metals are highly conductive. This is an important feature from the standpoint of corrosion because the metals provide an efficient path to transport electrons from the anode to the cathode. The resistivities of several metals at 20 °C (70 °C) are shown in Table 1. The microohm × cm unit (mW × cm) is 10–6 ohm × cm (W × cm); that is, a cube of material 1 by 1 by 1 cm has a resistance of 10–6 W from one face to the opposite face. All of these metals have high conductivities with respect to polymers or ceramics. Within the classification of metals, silver, copper, and gold have the highest conductivities. Iron is intermediate among the metals, and lead has one of the lowest conductivities. It should be noted that resistivity and conductivity are inverse relationships; that is, higher resistivities are equivalent to lower conductivities. Effect of Heat Treatment. Many mechanical properties of materials are improved by heat treatments. Unfortunately, such properties as hardness and strength are often achieved at the expense of corrosion resistance. For example, the hardness and strength of martensitic steels are counterbalanced by a lower corrosion resistance than for the ferritic and austenitic steels. The very high strengths achieved for precipitation-hardened steels are due to the secondary precipitates formed during the solution heat treating and aging process. As discussed above, precipitates with electrochemical properties distinctly different from those of the matrix have a deleterious effect on corrosion. Effect of Cold Working. Processes such as cold working, in which material is plastically deformed into some desired shape, lead to the formation of elongated and highly deformed grains and a decrease in corTable 1
Electrical resistivities of metals Temperature
Metal
Aluminum Brass Bronze Cadmium Chromium Cobalt Copper Gold Iron Lead Nickel Platinum Silver Zinc
°C
20 100 20 20 20 28 20 20 100 20 100 20 100 20 100 20 100 20 100 18 100 0 92.5
°F
70 212 70 70 70 82 70 70 212 70 212 70 212 70 212 70 212 70 212 64 212 32 199
Resistivity, mW × cm
2.828 3.86 7 18 7.6 13 9.8 1.724 2.28 2.44 2.97 10 16.61 22 27.8 6.141 10.327 10 14.1 1.629 2.15 5.76 8
Basic Concepts Important to Corrosion
rosion resistance. Cold working can also introduce residual stresses that make the material susceptible to stress-corrosion cracking. An improvement in corrosion resistance can be achieved by subsequently annealing at temperature at which grain recrystallization can occur. A partial anneal leads to stress relief without a major effect on the overall strength of the material. Effect of Welding. From the corrosion viewpoint, welding is a particularly troublesome treatment. Because welding involves the local heating of a material, it can lead to phase transformations and the formation of secondary precipitates. It can also induce stress in and around the weld. Such changes can lead to significant local differences in electrochemical properties as well as the onset of such processes as intergranular corrosion. Therefore, the weld filler metal should be as close in electrochemical properties to the base metal as technically feasible, and the weld should be subsequently stress relieved.
Inherent Reactivity Each metal has its own inherent tendency to corrode. Some metals, such as gold and silver, are very noble and have little tendency to corrode. They can be found in the earth in their natural, metallic state. At the other end of the scale of inherent reactivity are metals such as sodium. Sodium is an extremely active metal and corrodes spontaneously in the presence of water with a violent reaction. Iron is a moderately active metal and corrodes spontaneously in the presence of water. The natural state of iron in the structure of the earth is iron oxide. In order to recover iron from the iron oxide, a considerable amount of energy must be used to decompose the iron oxide and recover the pure iron. The variety of metals available provides a wide range of inherent reactivity, from very noble materials, which do not corrode readily, to extremely active metals, which corrode quite readily. An alternative way to express the inherent reactivity is to look at the amount of energy required to recover a metal from its oxide. Here, more energy is required for the most active metals; that is, it takes more energy to recover sodium from sodium oxide that it does to recover iron from iron oxide. Similarly, the noble metals have little tendency to form their oxides and are easily recovered from a metal oxide. The electromotive force (emf) series is a formal ranking of metals with respect to their inherent reactivity. Table 2 is an electromotive force series for many of the metals. The most noble metals are at the top of the emf series and have the highest positive standard electrode potentials. The most active metals are at the bottom of the series and have the most negative standard electrode potentials. The potential for hydrogen is taken as zero by internationally accepted convention. All other standard electrode potentials are referred to this standard hydrogen electrode
35
36
Corrosion: Understanding the Basics
(SHE) value. Thus, the potential of gold is +1.50 V with respect to the hydrogen reference potential, and the potential of iron is –0.44 V with respect to the hydrogen reference potential. The standard electrode potential values are determined for a special set of conditions; that is, the standard potential is for the equilibrium of the pure metal with its own ions at a specified concentration. No other ions are considered in the equilibrium. Thermodynamic calculations yield the standard electrode potential for each metal under these specified conditions. The emf series is most valuable for indicating the inherent reactivity of metals. Most corrosion applications, however, deal with mixed reactions, that is, not only the reaction of the metal with its own ions but also the reaction of the metal with other species in the solution, such as hydrogen ions or oxygen. In Chapters 3 and 4, various galvanic series for metals are discussed. These galvanic series take into account the other reactions and provide a listing of the inherent reactivity of metals in a specific environment. Nevertheless, the emf series is quite useful. Metals at the bottom or most negative end of the emf series are active metals. They have less corrosion resistance than metals higher in the series. In order of increasing corrosion resistance, magnesium is the least corrosion resistant, followed by zinc, iron, and copper, with gold being Table 2
Electromotive force series
Electrode reaction
Au3+ + 3e– ® Au Pd2+ + 2e– ® Pd Hg 2+ +2e– ® Hg Ag+ + e– ® Ag – Hg 2+ 2 + e ® 2Hg Cu+ + e– ® Cu Cu2+ + 2e– ® Cu 2H+ + 2e– ® H2 Pb2+ + 2e– ® Pb Sn + 2e– ® Sn Ni2+ + 2e– ® Ni Co2+ + 2e– ® Co Tl+ + 2e– ® Tl In3+ + 3e– ® In Cd2+ + 2e– ® Cd Fe2+ + 2e– ® Fe Ga3+ + 3e– ® Ga Cr3+ + 3e– ® Cr Cr2+ + 2e– ® Cr Zn2+ + 2e– ® Zn Mn2+ + 2e– ® Mn Zr4+ + 4e– ® Zr Ti2+ + 2e– ® Ti Al3+ + 3e– ® Al Hf4+ + 4e– ® Hf U3+ + 3e– ® U Be2+ + 2e– ® Be Mg2+ + 2e– ® Mg Na+ + e– ® Na Ca2+ + 2e– ® Ca K+ + e– ® K Li+ + e– ® Li
Standard potential at 25 °C (77 °F), V-SHE
1.50 0.987 0.854 0.800 0.789 0.521 0.337 0.000 (Reference) –0.126 –0.136 –0.250 –0.277 –0.336 –0.342 –0.403 –0.440 –0.53 –0.74 –0.91 –0.763 –1.18 –1.53 –1.63 –1.66 –1.70 –1.80 –1.85 –2.37 –2.71 –2.87 –2.93 –3.05
Basic Concepts Important to Corrosion
the most corrosion resistant. A metal with a more negative potential in the series will replace from solution the metal ions of a metal more positive in the series. Iron immersed in a solution containing copper ions, for example, a solution of copper sulfate, replaces the copper ions. The iron goes into solution as iron ions, and the copper ions plate out of solution onto metal as metallic copper. The position of metals in the emf series tells, in general, of the reactivity of the metal with deaerated acids, that is, acids containing no dissolved oxygen. Metals more negative than the hydrogen potential react with deaerated acids. Metals more positive than the hydrogen electrode are not attacked by deaerated acids. This only applies to acids in the absence of oxygen. The emf series provides an indication of the potential difference between two metals coupled together in a galvanic corrosion cell. If two dissimilar metals are coupled together, the potential difference is a driving force for corrosion reactions. The farther apart the metals are in the emf series, the greater the driving force for corrosion. For example, copper and aluminum form a strong cell with a potential difference of greater than 2 V, while magnesium and aluminum form a weaker cell with a potential difference of less than 1 V. The metal that is more positive in the series is the cathode, and the metal that is more negative in the series is the anode. The anodic member of the galvanic couple is severely corroded. Other considerations for the galvaniccorrosion couple exist, and the emf series should only be used as a general indication. These other considerations are discussed in subsequent chapters. In summary, the emf series ranks the metals with respect to their tendency to react. The metals at the top of the series are the most corrosion resistant and have the least tendency to be oxidized. An alternative way of expressing this is that the oxidizing power of a solution must be greater in order to corrode a metal higher in the series. It takes a solution with only a low oxidizing potential to corrode the active members of the series, such as sodium, magnesium, and aluminum. A greater oxidizing potential is required to corrode iron and nickel. An even greater oxidizing power of the environment is required to corrode copper, and a very high oxidizing power of the solution is required to corrode platinum and gold. This inherent reactivity of the metals is an important consideration in corrosion.
Formation of Corrosion Products The term corrosion products refers to the substances produced during a corrosion reaction. Corrosion products can be soluble, such as zinc chloride, which is formed when zinc is placed into a dilute hydrochloric acid, or zinc sulfate, which is formed when zinc is placed in
37
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Corrosion: Understanding the Basics
sulfuric acid, or corrosion products can be insoluble compounds, such as iron oxide or hydroxide. The presence of corrosion products is one way corrosion can sometimes be detected (e.g., rust). The tendency to form insoluble corrosion products is central to corrosion considerations because it is often the insoluble corrosion product films that provide passivity. The insoluble corrosion product film formed on a surface can block the surface from further attack and thus significantly reduce the corrosion rate of a material. The typical compounds of interest are oxides, hydroxides, and sulfates. Lead in sulfuric acid of specific concentrations, for example, forms a protective lead sulfate layer, and the corrosion rate is greatly reduced. Aluminum exposed to air develops a very tenacious and protective aluminum oxide passive film. Referring back to the emf series, aluminum is an extremely reactive metal. However, because of the formation of a very adherent and protective oxide layer, aluminum can be used for many architectural and structural purposes. Iron alloyed with a minimum of 12% Cr spontaneously forms a protective passive film on its surface. This spontaneous passivity for alloys with high chromium concentrations is the basis for an entire series of stainless steel alloys. The tendency to form an insoluble product is expressed as a solubility product. The solubility product defines the concentration of dissolved metal ions and, for example, hydroxyl ions required for the precipitation of a metal hydroxide. As the metal-ion concentration increases and the hydroxyl-ion concentration increases, the likelihood of formation of an insoluble product increases. Materials such as ferric hydroxide have extremely low solubility products, and it takes only a very small amount of ferric ion in solution to lead to the precipitation of ferric hydroxide. From the perspective of corrosion control, it is important to understand which products are the most stable and what degree of protection is provided by the solid products. Some insoluble products, such as the aluminum oxide that protects aluminum from corrosion, are very adherent to the metal surface and very dense. Other insoluble products are less dense and can be porous, and therefore provide little or no protection from subsequent corrosion.
Important Solution Characteristics In the previous section, the important characteristics of metals with respect to corrosion were discussed. In this section, the important characteristics of aqueous solutions are presented. These characteristics include: conductivity of the solution, acidity and alkalinity, oxidizing power, degree of ionization, and solubility in the solution. These characteristics of the solution, in combination with the characteristics of the
Basic Concepts Important to Corrosion
39
metal, will determine the corrosion behavior of a metal/environment combination. The conductivity of a solution is a measure of its ability to transport current. A high-conductivity solution easily transports current, whereas a solution with low conductivity transports current much less effectively. The conductivity is inversely proportional to resistivity; that is, if conductivity increases, resistivity decreases, and, conversely, if conductivity decreases, resistivity increases. Various solutions exhibit a wide range of conductivites. Seawater is a highly conductive solution and has a very low resistance to transporting current. Distilled water, on the other hand, is a very low-conductivity solution and has a high resistance to the transport of current. In general, as the concentration of dissolved species in the solution increases, the conductivity increases, and, in general, as the conductivity of the solution increases, the corrosion of metals in that solution increases. Recall that one of the requirements for a corrosion cell is an ionic conducting path between the anode and cathode. If the resistance of that ionic path is lower, there is less resistance in the corrosion cell, and the corrosion rate can proceed more rapidly. Seawater is much more corrosive than distilled water in part because of its much greater conductivity. The resistivity and conductivity of a one normal (1N) solution of several chemical compounds is presented in Table 3. The compounds are listed with their chemical formulas. The resistivity in ohm × cm and the conductivity in mho × cm are presented. Note that the conductivity values are simply the inverse of the resistivity; that is, to get conductivity, divide 1 by the value of the resistivity. Sodium chloride has a resistivity of 11.6 W × cm. This is a very low resistivity for a solution. It can be compared to the resistance of greater than 106 W for distilled water. Most of the solutions listed in Table 3 are reasonably low-resistance or Table 3
Resistivity and conductivity of 1 N solutions at 20 °C (70 °F) Resistivity, W · cm
Compound and formula
Boric acid, H3BO3(a) Chromic acid, H2CrO4 Copper sulfate, CuSO4 Ferrous chloride, FeCl2 Hydrochloric acid, HCl Nickel chloride, NiCl2 Nickel sulfate, NiSO4 Potassium chloride, KCl Potassium cyanide, KCN Potassium hydroxide, KOH Sodium chloride, NaCl Sodium hydroxide, NaOH Sodium carbonate, Na2CO3 Sodium sulfate, Na2SO4 Sulfuric acid, H2SO4 Zinc sulfate, ZnSO4
70,000 3.18 3.41 16.5 3.01 14.1 33.8 8.94 8.21 5.07 11.6 5.77 19.1 16.8 4.81 33.2
(a) Concentration is one molar (1M) instead of one normal (1 N).
Conductivity, mho · cm
0.000014 0.314 0.293 0.0606 0.3322 0.0709 0.0295 0.1119 0.1218 0.1972 0.0862 0.1733 0.0523 0.0595 0.208 0.0301
40
Corrosion: Understanding the Basics
high-conductivity solutions. The exception is boric acid, which is a moderately high-resistance medium. Compare the values of sulfuric acid and hydrochloric acid, which are quite low in resistivity, with that of boric acid, which is quite high. All other things being equal, boric acid would be considerably less corrosive than sulfuric or hydrochloric acids at similar concentrations. The relative acidity or alkalinity of a solution greatly affects its corrosivity for particular metals. Solutions can be described as acidic, neutral, or alkaline based on the relative ratio of hydrogen ions to hydroxyl ions. Figure 14 shows this relationship. Where hydrogen ions and hydroxyl ions are in balance, the solution is neutral. Where hydrogen ions predominate over hydroxyl ions, the solution is acidic, and where the hydroxyl ions predominate over the hydrogen ions, the solution is alkaline. Strongly acidic solutions have a greater number of hydrogen ions, and strongly alkaline solutions have a greater concentration of hydroxyl ions. The pH defines the acidity or alkalinity of a solution. The pH is defined as –log(H+); an increase of one pH unit is equivalent to an order of magnitude, or factor of 10, decrease in the hydrogen ion concentration. Figure 15 shows a pH scale ranging from 1 to 14. A value of pH 7 defines a neutral solution, and low values of pH identify the solution as being acidic. The lower the value, the stronger the acid becomes. High pH values, greater than 7, identify the solution as an alkaline solution, and the higher the pH, the stronger the alkaline environment becomes. Also shown in Figure 15 are some common environments and their typical positions along the pH scale. Hydrochloric, sulfuric, and nitric acids are strong acids and have low pH values even in relatively dilute solutions. Boric, citric, and phosphoric acids are weaker acids and have pH values only slightly acidic. Tap water and seawater typically have neutral pH values. Sodium bicarbonate and ammonium hydroxide are mildly alka-
Fig. 14
Range of acidity and alkalinity
Fig. 15
The pH of several common environments
Basic Concepts Important to Corrosion
41
line solutions and have a pH of approximately 10. Sodium hydroxide is a strongly alkaline solution and has high pH values. The approximate pH of solutions of acids and bases is shown in Table 4. The strong acids have the lowest values of pH, and the strong bases have the highest values of pH. The pH of an acid becomes lower as the concentration increases; that is, the acid becomes stronger. The pH of a strong base becomes greater as the concentration increases; that is, the solution becomes more highly alkaline. The effect of acidity or alkalinity of the solution depends very much on the specific metal of interest. For example, nickel is quite resistant to highly alkaline environments, whereas aluminum is severely corroded by strongly alkaline environments. Again, it is important to discuss only the corrosion behavior of a combination of a metal in a specific environment. The oxidizing power of a solution is a measure of its relative tendency to corrode or oxidize metals. A solution of low oxidizing power corrodes only those metals at the lower (more active) end of the electromotive force series. A solution of strong oxidizing power corrodes all metals on the series except those with the most positive (most noble) values of the emf series. The range of oxidizing power encountered in aqueous environments is from strongly oxidizing to strongly reducing. The oxidizing power is an inherent property of the chemical species. Increasing oxidizing power means the tendency for the solution to oxidize a metal increases. Figure 16 compares the tendency of a metal to corrode, as expressed by the emf series, with the oxidizing power of the solution, ranging from highly oxidizing at the top to highly reducing at the bottom. Magnesium has the greatest tendency to corrode, and gold Table 4
Approximate pH of solutions of acids and bases
Solution
Normality, N
pH
Acids Hydrochloric acid Sulfuric acid Orthophosphoric acid Formic Acetic Boric acid
1 0.1 0.01 1 0.1 0.01 0.1 0.1 1 0.1 0.01 0.1
0.1 1.1 2 0.3 1.2 2.1 1.5 2.3 2.4 2.9 3.4 5.2
1 1 0.1 0.01 0.1 0.1 1 0.1 0.01 0.1
14 14 13 12 12 11.6 11.6 11.1 10.6 11
Bases Potassium hydroxide Sodium hydroxide Sodium and potassium hydroxide Trisodium phosphate Sodium carbonate Ammonium hydroxide Potassium cyanide
42
Corrosion: Understanding the Basics
has the least tendency to corrode. Concentrated nitric acid is a highly oxidizing environment, aerated (containing oxygen) acid is mildly oxidizing, and deaerated (no oxygen) acid is a relative reducing environment. The oxidizing power of a deaerated acid is sufficient to corrode both magnesium and iron but is insufficient to corrode copper or gold. An aerated acid, that is, one containing dissolved oxygen, has sufficient oxidizing power to corrode magnesium, iron, and copper. An aerated acid is still insufficient to corrode gold. Concentrated nitric acid has a highly oxidizing power and corrodes gold, copper, iron, and magnesium. The addition of oxygen dissolved in a solution increases its oxidizing power. Other chemical species, however, also increase the oxidizing power. Ferric ions and cupric ions greatly increase the oxidizing power of the solution. As mentioned previously, concentrated nitric acid is a strongly oxidizing environment. Deaerated hydrochloric acid is an example of a highly reducing environment. Dissociation or Ionization. The corrosivity of an environment is strongly dependent upon the degree of ionization of chemical species in the solution. Ionization or dissociation is the separation of the chemical into ionic species, examples of which include the following: sodium chloride, NaCl ® Na+ + Cl–; sulfuric acid H2SO4 ® H+ + HSO -4 and HSO -4 ®H+ + SO -4 ; and sodium hydroxide, NaOH®Na+ + OH–.. Sodium chloride (NaCl) dissolved in water ionizes to form a sodium ion (Na+) and a chloride ion (Cl–). Sulfuric acid can ionize to form a hydrogen ion plus a negatively charged species and can further dissociate to form a second hydrogen ion plus a sulfate ion. Sodium hydroxide ionizes to form a sodium ion and a hydroxyl ion. The degree of ionization or the number of sodium chloride molecules that break up into sodium ions and chloride ions depends on the particular compound and its concentration.
Fig. 16
Relationship between the tendency of a metal to corrode and the oxidizing power of a solution
Basic Concepts Important to Corrosion
43
The degree of ionization of several chemicals is shown in Table 5. Complete ionization is represented by a degree of ionization equal to 1.0. Values close to 1.0 for the degree of ionization indicate that the compound forms many ions, and very low values indicate that the compound forms only a few ions. Hydrochloric acid and potassium hydroxide have high degrees of ionization, and these compounds dissociate to a large degree, forming ions in solution. Boric acid and ammonium hydroxide have low degrees of ionization and form fewer ions in solution. This accounts for hydrochloric acid and potassium hydroxide being a strong acid and a strong base, respectively. These compounds also produce a large number of hydrogen ions (hydrochloric acid) and hydroxyl ions (potassium hydroxide) in solution. Boric acid and ammonium hydrogen, on the other hand, are a weak acid and a weak base, respectively, because there are only a few hydrogen ions and hydroxyl ions in their solutions. The chemicals that are acids generate hydrogen ions when they dissociate. The chemicals that are bases produce hydroxyl ions when they dissociate. Salts are chemicals that produce neither hydrogen ions nor hydroxyl ions when they dissociate. Hydrochloric acid produces a large number of hydrogen ions. Potassium hydroxide produces a large number of hydroxyl ions, and sodium chloride produces only sodium ions and chloride ions, that is, neither hydrogen ions nor hydroxyl ions. Solubility is a measure of the quantity of an ion or gas in a solution. There is a saturation limit, or upper limit, on solubility for species in a given solution. For example, if one starts with a solution containing no oxygen and bubbles oxygen through that solution, the oxygen dissolved in the solution begins to increase. The dissolved oxygen concentration continues to increase until the saturation limit is reached. At the saturation limit, the addition of more oxygen simply bubbles through the solution Table 5
Degree of ionization of acids, bases, and salts at 25 °C (77 °F)
Solution
Degree of ionization
[H+]
Acids Hydrochloric acid, 1 N Hydrochloric acid, 0.5 N Sulfuric acid, 1 N Hydrofluoric acid, 1 N Boric acid, (primary ionization), 0.1 M Hydrocyanic acid, 0.1 M Phosphoric acid, (secondary ionization), 0.5 N
0.784 0.876 0.510 0.070 0.0001 0.0001 0.170
–
Bases Potassium hydroxide, 1 N Sodium hydroxide, 1 N Ammonium hydroxide, 1 N
[OH ] 0.77 0.73 0.004
Salts Such as KCl, 0.1 N Such as K2SO4, Na2AO4, 0.1 N Such as CuSO4, NiSO4, 0.1 N
0.784 N 0.438 N 0.510 N 0.070 N 0.00001 N 0.00001 N 0.085 N
0.77 N 0.73 N 0.004 N [Metal ion]
0.86 0.72 0.45
0.86 N 0.072 N 0.045 N
44
Corrosion: Understanding the Basics
without increasing the amount dissolved. Similarly, solids dissolve in a solution and continue to increase in concentration until their saturation limit is reached. Beyond the saturation limit, solid deposits form in a solution and precipitate from solution. Some of the solids that form can provide a protective film (passive film) and reduce the corrosion rate of the metal. Complexing agents can combine with ions in solution and increase the apparent solubility of those ions. This is done by tying up a number of those ions in the form of soluble complexes. A result of the presence of complexing agents in a solution can be the prevention of the formation of a protective film. An example of a complexing agent is ammonia with copper. The presence of ammonia species in the solution greatly increases the solubility of copper ions and consequently increases the corrosion rate of copper. The formation of protective, insoluble products on the copper surface is retarded by the complexing species. From a corrosion perspective, the solubility of oxygen in a solution is one of the most significant effects. Figure 17 shows the effect of increasing oxygen concentration on the corrosion rate of iron in water. At any given temperature, the corrosion rate of iron increases with increasing oxygen concentration. This figure also shows the increase in the corrosion rate of iron with increasing temperature at any given oxygen concentration. Table 6 provides corrosion rates for various metals
Fig. 17
The effect of oxygen concentration on the corrosion rate of iron
Table 6 Comparison of corrosion rates (in mm/yr) in oxygen-free (hydrogen-saturated) and oxygen-saturated solutions Metal
Mild steel Lead Copper Tin Nickel Monel (a) No oxygen
Acid
Hydrogen saturated(a)
6% H2SO4 4% HCl 4% HCl 6% H2SO4 4% HCl 2% H2SO4
40 35 25 9 9 2
Corrosion rate Oxygen saturated
500 325 2150 1090 675 140
Basic Concepts Important to Corrosion
Table 7
Oxygen solubility in water Temperature
°C
0 20 40 60 80 100
45
°F
32 70 105 140 140 212
Oxygen solubility Grams per kg water Parts per million
0.069 0.043 0.031 0.027 0.014 0.00
69 43 31 27 14 <1
in solutions with no oxygen (hydrogen-saturated) and in solutions saturated with oxygen. For all of the metals listed, the addition of oxygen to the solution significantly increases the corrosion rate. The presence of oxygen in a solution provides an additional reduction process, that is, an additional reaction to consume electrons. As stated earlier in this chapter, iron atoms going into solution leave electrons at the anodic sites, and the corrosion reaction can only go on as quickly as these electrons can be consumed in a cathodic reaction. Thus, the addition of oxygen allows more electrons to be consumed by an additional reduction reaction, and the corrosion rate of many metals increases. The oxygen solubility in water as a function of temperature is shown in Table 7. The oxygen solubility decreases as temperature increases from 0 °C (32 °F) through 100 °C (212 °F). At the boiling point, all oxygen is stripped from the water, and the solubility essentially becomes zero. The corrosion rate of an oxygen-saturated solution based on oxygen solubility would be predicted to decrease with increasing temperature. This effect is often offset by increasing reaction kinetics as temperature increases; however, the corrosion rate drops rapidly at the boiling point because of a discontinuous drop in oxygen concentration.
Corrosion Rate Expressions and Allowances
Corrosion Rate Expressions. The term corrosion rate refers to the effect of general (or uniform) corrosion on a metal (or nonmetal) per unit of time. The type of corrosion rate expression used depends on the technical system and on the type of corrosion effect. Thus, corrosion rate can be expressed as (a) the mass of metal turned into corrosion products per unit area of surface per unit of time (weight loss), or (b) an increase in corrosion depth per unit time (penetration rate). The corrosion effect, which refers to a change in any part of the corrosion system (metal/environment combination), can vary with time and might not be the same at all points of the corroding surface. Reports of corrosion rates, therefore, should be accompanied by information on the type, time dependency, and location of the corrosion effect.
46
Corrosion: Understanding the Basics
Weight-loss tests are the most common of all corrosion-test measurements for determining corrosion rates. A clean coupon is measured, weighed, exposed to the corrodent for a known time, removed, cleaned to remove corrosion products, and reweighed. The rate of metal removal due to corrosion is then calculated from: R = KW/ATd where R is the corrosion rate, K is a constant, W is weight loss in g to the nearest mg (corrected for any loss during cleaning), A is the area to the nearest 0.01 cm2, T is time of exposure in hours to the nearest 0.01 h, and d is the density in g/cm3. Many different units are used to express the corrosion rate, R. Using the preceding units for W, A, T, and d, the corrosion rate can be calculated in a variety of units with the appropriate value of K: Corrosion rate units desired (mpy)
Constant (K) in corrosion rate equation
3.45 × 106 3.45 × 103 2.87 × 102 8.76 × 104 8.76 × 107 2.78 × 106 1.00 × 104 × d(a) 2.40 × 106 × d(a) 2.78 × 106 × d(a)
Mils per year (mpy) Inches per year (in./yr) Inches per month (ipm) Millimeters per year (mm/yr) Micrometers per year (mm/yr) Picometers per second (pm/s) Grams per square meter per hour (g/m2/h) Milligrams per square decimeter per day (mdd) Micrograms per square meter per second (mg/m2/s)
(a) Density is not needed to calculate the corrosion rate in these units. The density in the constant K cancels out the density in the corrosion rate equation.
If desired, these constants can also be used to convert corrosion rates from one set of units to another. To convert corrosion rate in units, X, to a rate in units, Y, multiply by KY/KX. For example, if R is 15 mpy, the rate in pm/s would be calculated from the following: æ 2.78 ´ 10 6 15ç ç 3.45 ´ 10 6 è
ö ÷ = 12.1 pm/ s ÷ ø
In the United States, the most widely used expression for quantitatively defining the rate of corrosion in mils per year (abbreviated mpy or mils/yr) and calculated from: Mils per year = 534W/dAT where W is the weight loss in mg; d is the density, g/cm3; A is the specimen area, in 2. ; and T is the exposure time, h. Throughout this book, metals and nonmetals are compared on the basis of their corrosion resistance measured in mpy and the appropriate metric equivalent.
Basic Concepts Important to Corrosion
47
Relationships among some of the units commonly used for corrosion rates are given in Table 8. When to Use Metals and Alloys in Certain Corrosion Rates. As discussed in later chapters, different metals and alloys have varying corrosion rates depending, of course, on specific environments and environmental variables. Table 9 lists some general guidelines, which can serve as a starting point when trying to determine maximum corrosion rates for specific working conditions for cases of general (uniform) corrosion. Corrosion Allowance. By knowing the expected general corrosion rate and the anticipated plant or service life of a part, the designer can calculate the extra wall thickness required for corrosion resistance of the process equipment being designed. After determining a wall thickness that meets mechanical requirements, such as pressure and weight of equipment, an extra thickness called a corrosion allowance is added to the wall thickness to compensate for the metal expected to be lost over the life of the equipment. Then, because the penetration depth can vary, a corrosion allowance is assigned a safety factor of 2. As an example, suppose a tank wall required a 5 mm (316 in.) wall thickness for mechanical considerations. The designer has determined that the corrosion rate will be 15 mpy (0.4 mm/yr) and the expected life of the equipment will be 10 yr. The total corrosion allowance is the corrosion rate per year (0.4 mm, or 0.015 in.) × 10 yr = 4 mm (0.15 in.). The corrosion allowance is doubled to 0.3 in. (8 mm) as a safety consideration. Table 8
Relationships among some units commonly used for corrosion rates
d is metal density in grams per cubic centimeter (g/cm3). Unit
mdd
g/m2/d
Factor for conversion to mm/yr mm/yr
Milligrams per square decimeter 1 0.1 36.5/d per day (mdd) 10 1 365/d Grams per square meter per day (g/m2/d) Microns per year (mm/yr) 0.0274d 0.00274d 1 Millimeters per year (mm/yr) 27.4d 2.74d 1,000 Mils per year (mils/yr) 0.696d 0.0696d 25.4 Inches per year (in./yr) 696d 69.6d 25,400
Table 9
10 max 20 max 50 max Rates > 50
0.365/d 0.001 1 0.0254 25.4
mils/yr
1.144/d
in./yr
0.00144/d
14.4/d
0.0144/d
0.0394 39.4 1 1,000
0.0000394 0.0394 0.001 1
Characteristics and uses of corrosion rates
Penetration rate, mpy
1 max
0.0365/d
Characteristics and uses
Very low corrosion; recommended for services where product contamination is a problem, e.g., food industry equipment Low corrosion; recommended for thin-walled process equipment Fairly low corrosion; can be considered the normal maximum allowed in chemical equipment High corrosion; seldom tolerated except in thick-walled equipment where product contamination is controlled Excessive corrosion; very seldom tolerated and only then in very thick-walled equipment where massive product contamination is not a problem
48
Corrosion: Understanding the Basics
The final wall thickness would be 0.3 + 0.1875 = 0.4875 in. (8 + 5 = 13 mm). The designer would then specify a 1 2 in. (13 mm) wall thickness as the closest standard plate available. Additional information on corrosion allowance calculations can be found in Chapter 7.
References Selected References Corrosion · Corrosion, Vol 13, ASM Handbook, ASM International, 1987 · Corrosion Basics: An Introduction, L.S. Van Delinder, Ed., National Association of Corrosion Engineers, 1984 · M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill, 1986
General Metallurgy · Metals Handbook Desk Edition, 2nd ed., J.R. Davis, Ed., ASM International, 1998 · Metallurgy for the Non-Metallurgist, H. Chandler, Ed., ASM International, 1998
Corrosion: Understanding the Basics J.R. Davis, editor, p49-97 DOI: 10.1361/cutb2000p049
CHAPTER
Copyright © 2000 ASM International® All rights reserved. www.asminternational.org
3
Principles of Aqueous Corrosion CORROSION OF METALS in aqueous environments is electrochemical in nature involving two or more electrochemical reactions taking place on the metal surface. As a result, some of the elements of the metal or alloy change from a metallic state into a non-metallic state. The products of corrosion can be dissolved species or solid corrosion products; in either case, the energy of the system is lowered as the metal converts to a lower-energy form. Rusting of steel is the best known example of conversion of a metal (iron) into a nonmetallic corrosion product (rust). The change in the energy of the system is the driving force for the corrosion process, which behaves according to the laws of thermo-dynamics. The thermodynamics of aqueous corrosion is the subject of the first half of this Chapter. Important concepts described include the following: · · · ·
Corrosion reactions and free-energy change The relationship between free energy and electrochemical potential The effect of ionic concentration on electrode potential The corrosion behavior of a metal based on its potential-pH diagram
As indicated above, corrosion is an electrochemical process. Electrochemical processes require anodes and cathodes in electrical contact and an ionic conduction path through an electrolyte (see, for example, Fig. 2 in Chapter 2). The electrochemical process includes electron flow between the anodic and cathodic areas; the rate of this flow corresponds to the rates of the oxidation and reduction reactions that occur at the surfaces.
50
Corrosion: Understanding the Basics
Understanding the kinetics of corrosion and the factors that control the rates of corrosion reactions requires examination of the concepts of polarization behavior and identification of the various forms of polarization in an electrochemical cell. As described in the second half of this chapter (see the section “The Kinetics of Aqueous Corrosion”), these concepts include anodic and cathodic reactions, the mixed-potential theory, and the exchange currents.
The Thermodynamics of Aqueous Corrosion Thermodynamics is a powerful science that defines which reactions are possible and whether a particular reaction will occur. It provides a sound basis for the understanding of corrosion phenomena and is central to the study of corrosion processes. This section identifies the important concepts of thermodynamics relevant to corrosion and discusses applications of those concepts. Thermodynamics describes equilibria as a function of the elements and compounds present and the environmental conditions, such as pressure, temperature, and chemical composition. Thermodynamics is used to determine whether corrosion can occur and to predict the stable corrosion products that will form. Thermodynamic concepts are referenced throughout this book to explain observed corrosion behavior. For example, copper will not corrode in a strong hydrochloric acid (HCl) solution if oxygen is not present; if oxygen is dissolved in the HCl, however, copper corrodes at a rapid rate. This behavior is readily explained using thermodynamic concepts.
Corrosion Reactions and Free-Energy Change A law of nature is that the most stable state for a set of reactants is that state which has the lowest free energy. Consequently, a metal in contact with a solution moves toward the lowest free-energy state. When the system arrives at this state, there is no further change. This final, unique, lowest-energy state is called equilibrium. At equilibrium, the system is stable, and there is no driving force for any change from that state. This concept can be demonstrated by considering the corrosion of iron in water. When iron is immersed in water, the corrosion reaction of interest is the reaction between iron atoms in the metal and the corrosion products of the iron that is, ferrous ions (Fe2+) in solution. This reaction can be expressed as the following: Reacts ¾® Iron atom in metal Ferrousion in solution ¬ ¾
or
Principles of Aqueous Corrosion
51
Fe 2 + + 2e – « Fe The reaction shows the equilibrium between dissolved ferrous ions in solution and metallic iron. The reaction can move either from left to right or from right to left. The first type of reaction results in ferrous ions combining with electrons to form new iron atoms at the metal surface. This is equivalent to a metal plating process. The reaction from right to left involves the removal of an iron atom from the metal surface with the formation of a new ferrous ion (Fe2+) in solution. Two electrons are left at the iron surface for each iron atom that becomes a ferrous ion. This type of reaction is equivalent to corrosion and results in metal loss. When at equilibrium, the reaction proceeds at the same rate in both directions, and there is no net change in ferrous ion concentration or weight loss of the iron. If the free energy of the system is lowered by a decrease in the ferrous ion concentration, then there is a greater driving force for the reaction to proceed from left to right, with a net reduction in ferrous ion concentration. As the reaction proceeds, the driving force is lowered, and the system moves toward its equilibrium state. If, on the other hand, the system would have a lower free energy by increasing the concentration of ferrous ions in solution, there would be a driving force for the reaction to proceed from right to left. As this occurs, iron atoms at the metal surface react to form ferrous ions in solution. This reaction of iron atoms increases the ferrous ions in solution. Also, the driving force for further reaction would be decreased, and the system would move toward its equilibrium. As stated previously, if the system is at equilibrium, there is no driving force for a net reaction from either left to right or right to left, and the ferrous ion concentration remains constant. This latter case is the specific concentration that results in equilibrium. This phenomenon can be illustrated with the use of free-energy diagrams. Three diagrams are shown in Fig. 1; each represents one of the cases described above. On the diagrams, the free-energy change is identified as DG. The free energy of a ferrous ion in solution is represented
Fig. 1
Free-energy diagrams for reactions of ferrous ions and iron atoms
52
Corrosion: Understanding the Basics
by the bottom of the trough under the Fe2+ symbol, and the free energy of an iron atom in the metal surface is shown by the bottom of the trough beneath the Fe symbol. In all cases, the reaction of a ferrous ion in solution to become an iron atom on the metal surface is considered. The net free-energy change for the reaction is shown by comparing the relative levels of the bottom of each trough. For the diagram on the left of Fig. 1, the trough for the ferrous ion in solution is more positive than the trough for the iron atom. This reaction proceeding from left to right results in a net decrease in free energy; that is DG is less than zero. For this case, the system could lower its free energy by the reaction of ferrous ions to form metal atoms on the surface; the plating process will occur. The reaction will proceed from left to right. For the case shown in the diagram on the right of Fig. 1, the trough for the ferrous ion is more negative than the trough for the iron atom. The reaction of ferrous ions to form metal atoms will result in an increase in free energy; that is, DG is positive. This reaction as written will not proceed, because it involves an increase in free energy; rather, the reaction proceeds in the reverse direction. Iron atoms go into solution as ferrous ions, thus reducing the free energy of the system. The diagram in the center of Fig. 1 represents the equilibrium condition. In this case, the trough for the ferrous ion is at the same level as the trough for the iron atom. Under these conditions, there is no change in free energy for the reaction from either right to left or left to right; that is, DG is equal to zero. For this case, there is no driving force for either the plating process or the corrosion process, and the system remains at the equilibrium state. The ferrous ion concentration in solution remains constant, and there is no corrosion of the iron in this solution. An iron sample exposed under these conditions would have the same weight after any length of exposure as it had on being immersed in the solution. Another behavioral characteristic is that any system will move toward equilibrium. The system will react in a manner to offset any driving force for reactions, and equilibrium is eventually obtained when there is no net driving force, or DG = 0, for the reaction in either direction. For the diagram on the left of Fig. 1, the ferrous ion concentration is greater than the equilibrium concentration. Because of this, there is a net driving force to consume some of those ferrous ions and to deposit metal atoms on the surface, thus reducing the ferrous ion concentration in solution. For the diagram on the right of Fig. 1, the ferrous ion concentration is lower than the equilibrium concentration. Because of this, there is a driving force for the production of ferrous ions by the removal of metal atoms from the surface and an increase in ferrous ion concentration, thus moving toward the equilibrium concentration. To summarize, a metal in a solution has a characteristic free-energy change with respect to metal atoms at the surface of the metal and metal ions in solution, which is determined by such factors as the composition
Principles of Aqueous Corrosion
53
of the metal, the chemical composition of the solution, temperature, and pressure. A law of nature is that systems react to minimize their free energy. The reactions between metal atoms at the metal surface and metal ions in solution proceeds to lower the free energy of the system. The reaction removes metal ions from solution and increases the metal atoms at the surface (metal plating) or proceeds to remove metal atoms from the surface and produces metal ions in solution (corrosion). The direction of the reaction depends on the relative change in free energy. When there is no difference in free energies between the metal atoms and the metal ions in solution, the system is at equilibrium and no further net reaction occurs. The preceding discussion focuses on the equilibrium behavior of ferrous ions and iron. A general relationship for any metal ion in solution and the corresponding metal atom on the surface is: Mn+ + ne– « M This general electrode reaction refers to the reaction of ions of the metal (M) with electrons to produce uncharged metal atoms of M on the metal surface. The charge of the metal ion and the number of electrons involved in the reaction are identified by n and is called “valence.” The value of n is a characteristic of the particular metal. Electrode reactions for several metals with their corresponding values of n are presented in Table 1. The characteristic value for gold is 3+ and the characteristic value for iron and zinc is 2+. The characteristic value for aluminum is 3+. Some metals have alternative valences, and multiple values of n are listed, for example, n values of 1+ or 2+ for copper.
Free Energy and Electrochemical Potential The relationship between free energy (DG) and electrochemical (or cell) potential (E) is described by the equation: DG = –nFE where n is the number of electrons in the reaction and F is Faraday’s constant. The free-energy change of the reaction is equal to the negative of the product of the number of the electrons in the reaction times a constant value (Faraday’s constant) times the electrode potential. Large negative free-energy changes give rise to large positive potential differences, and large positive free-energy changes give rise to large negative potential differences. These terms are equivalent in that they both describe the magnitude of the driving force for a reaction to occur. Furthermore, at equilibrium, where there is no driving force for the reaction, both the free-energy change (DG) and the driving force in terms of potential (E)
54
Corrosion: Understanding the Basics
Table 1 Electrode reactions of several metals and their ions Metal
Electrode reaction(a)
Aluminum Beryllium Cadmium Calcium Chromium
Al3+ + 3e– « Al Be2+ + 2e– « Be Cd2+ + 2e– « Cd Ca2+ + 2e– « Ca Cr3+ + 3e– « Cr Cr2+ + 2e– « Cr Co2+ + 2e– « Co Cu+ + e– « Cu Cu2+ + 2e– « Cu Ga3+ + 3e– « Ga Au3+ + 3e– « Au Hf4+ + 4e– « Hf 2H+ + 2e– « H2 In3+ + 3e– « In Fe2+ + 2e– « Fe Pb2+ + 2e– « Pb Li+ + e– « Li Mg2+ + 2e– « Mg Mn2+ + 2e– « Mn Hg2+ + 2e– « 2Hg – Hg 2+ 2 + 2e « 2Hg Ni2+ + 2e– « Ni Pd2+ + 2e– « Pd K+ + e– « K Ag+ + e– « Ag Na+ + e– « Na Tl+ + e– « Tl Sn2+ + 2e– « Sn Tl2+ + 2e– « Ti U3+ + 3e– « U Zn2+ + 2e– « Zn Zr4+ + 4e– « Zr
Cobalt Copper Gallium Gold Hafnium Hydrogen Indium Iron Lead Lithium Magnesium Manganese Mercury Nickel Palladium Potassium Silver Sodium Thallium Tin Titanium Uranium Zinc Zirconium
(a) Note the characteristic value for n, i.e., the number of electrons
are equal to zero. In order to maintain the signs appropriate to those conventions, two procedures are followed: · All reactions are written to consume electrons, e.g., for the reaction of ferrous ions plus electrons to produce uncharged metal atoms expressed as Fe2+ + 2e– ® Fe. · The relationship DG = –nFE is used.
Using these two procedures ensures that consistent values for potential and the proper sign for the driving force for a reaction are realized. The relationships among free energy and potential and the significance to the direction in which a reaction will proceed are shown in Table 2. For the reaction of metal ions plus electrons to produce metal atoms at the surface, the values of DG and E determine whether the reaction is at equilibrium or proceeds spontaneously from left to right or from right to left. When the driving force for the reaction, expressed either as DG or E, is equal to zero, the system is at equilibrium, and there is no further net reaction. When the free-energy change is less than zero or when the difference in electrode potential is greater than zero, the reaction proceeds spontaneously from left to right; that is, the system re-
Principles of Aqueous Corrosion
55
Table 2 Relationships among free energy, potential, and the direction in which the reaction (Mn+ + ne– « M) will proceed DG
0 <0 (negative) >0 (positive)
E
Spontaneous net reaction
0 >0 (positive) <0 (negative)
Equilibrium; no net reaction Spontaneous reaction; left to right; plating Spontaneous reaction; right to left; corrosion
acts with the consumption of metal ions and the production of metal atoms at the surface. When the free-energy change for the reaction is greater than zero or the difference in electrode potential is less than zero, the reaction proceeds spontaneously from right to left; that is, the reaction proceeds to remove metal atoms from the surface and produces metal ions in solution. This relationship between free-energy and the thermodynamics of electrochemical potential is central to the study of corrosion. Through the measurement and control of electrode potential, the free energy can be measured and controlled. A consequence of this is that an instrument that allows the change and control of potential to desired values is essentially a free-energy pump. The direction of reaction and the thermodynamically stable species can be controlled by the control of electrode potential.
Tendency for Metals to Corrode Each metal has a characteristic, inherent tendency to corrode or to react in an aqueous environment to produce metal ions. This inherent reactivity can be described in terms of either free-energy, DG, or electrochemical potential, E. The inherent reactivity of a metal can be expressed by the magnitude of the free-energy change on the metal going to a metal corrosion product. Table 3 presents the free-energy change for the reactions of four common metals to form their hydroxides. The first reaction is for magnesium reacting with water and oxygen to form a solid corrosion product, that is, magnesium hydroxide. The second reaction is for iron reacting with water and oxygen to form ferrous hydroxide, and the third is for copper reacting with water and oxygen to form cuprous hydroxide. The free-energy change, DG, measured in joules, is presented for each of these reactions. A negative value for the free-energy change indicates that the metal will react spontaneously to form the corrosion Table 3 Inherent reactivity of four metals expressed as the magnitude of free-energy change for a given corrosion reaction Metal
Magnesium Iron Copper Gold
Corrosion reaction
Mg + H2O(l) +
1
2 O2(g) « Mg(OH)2(s)
DG, kJ
–597
Fe + H2O(l) + 1 2 O2(g) « Fe(OH)2(s)
–249
« Cu(OH)2(s)
–120
O2(g) « Au(OH)3(s)
+66
Cu + H2O(l) + Au
+3
2
1
2 O2(g)
H2O(l) +
3
4
56
Corrosion: Understanding the Basics
product. It is important to note again that while thermodynamics indicates that a reaction will occur, it does not provide any information as to the rate of reactions. A positive free-energy change means that the reaction of the metal is not spontaneous and will not occur. The more negative the value of DG, the greater the inherent reactivity of the metal. A comparison of the free-energy change values for each reaction in Table 3 indicates that magnesium is the most reactive metal, followed by iron and copper. The positive value for the free-energy change of gold, whose corrosion reaction is also shown in Table 3, indicates that gold will not spontaneously react in the presence of water and oxygen. Gold is referred to as being noble or immune from attack under these conditions whereas magnesium, iron, and copper are active and will corrode.
Effect of Ionic Concentration on Electrode Potential Understanding the effect of ion concentration on electrode potential is facilitated by reference to the reaction Fe2+ + 2e– ® Fe, which shows that ferrous ions (Fe2+) plus two electrons (2e–) can be reduced to metallic iron (Fe). The Nernst equation, which is useful for determining the effect of ionic concentration on the electrode potential for such a reaction, is expressed as the following: RT ( red) (Eq 1) E = E0 ln nF (ox) where E is the electrode potential, E 0 is the standard electrode potential, R is the gas constant, T is the absolute temperature (in degrees Kelvin), n is the number of electrons involved in the reaction, F is the Faraday constant, and (red) and (ox) are the concentrations of reduced and oxidized species, respectively. The standard electrode potential is determined for the special condition of all elements in the reaction in a pure state and all ions at 1 M concentration. Electrode potential increases or decreases from the value of the standard electrode potential based on the ratio of concentration of the reduced species, in this case iron, and the concentration of the oxidized species, in this case the ferrous ion. For the ferrous ion reduction to iron reaction expressed as: 0.059 Fe (Eq 2) log 2 Fe 2 + the standard electrode potential is –0.440 V with respect to the standard hydrogen electrode (SHE). At 25 °C (75 °F), the value for the constant 2.3 times RT divided by F is 0.059. The 2.3 factor converts the natural log (ln) to log base 10 (log). The value for n in this equation is 2. The activity or concentration of a pure solid phase, for example, iron, is equal to 1. E = -0.440 -
Principles of Aqueous Corrosion
57
The final equation for the electrode potential as a function of concentration can be expressed as the following: E = –0.440 + 0.030 log Fe2+
(Eq 3)
This equation indicates that an order of magnitude, or factor of 10, increase in the ferrous ion concentration increases the electrode potential by 0.030 V. Conversely, a decrease in the ferrous ion concentration by a factor of 10 decreases the electrode potential by 0.030 V. The effect of concentration on the equilibrium electrode potential of two reactions is shown in Table 4. The ferrous ion reduction reaction is considered first. The standard electrode potential value is determined when pure iron is in equilibrium with ferrous ions at a 1 M concentration (unit activity). This is the value of –0.440 V-SHE. If the ferrous ion concentration is decreased by a factor of 1000, or three orders of magnitude, the resulting ferrous ion concentration is 10–3, and the electrode potential is decreased to –0.530 V-SHE. For an additional decrease in concentration by three orders of magnitude to 10–6, the corresponding electrode potential is –0.620 V-SHE, and for a concentration of 10–9 M ferrous ion concentration, the electrode potential is equal to –0.710 V-SHE. These values are calculated from Eq 1, 2, and 3. Similar calculations for the ferric ion reduction to ferrous ions yields the values shown at the bottom of Table 4. The reduction reaction involves the reaction of a ferric ion plus one electron being reduced to a Table 4 Effect of concentration on equilibrium electrode potential Ferrous ion reduction
Ferrous ion + 2 electrons ® iron Fe2+ + 2e– ® Fe æ ö E = –0. 440 – 0. 059 log ç 12+ ÷ 2 è Fe ø Ferrous ion, molar
1 10–3 10–6 10–9
E, V-SHE
–0.440 –0.530 –0.620 –0.710
Ferric ion reduction
Ferric ion + 1 electron ® ferrous ion Fe3+ + e– ® Fe2+ æ Fe 2 + E = 0. 771 – 0. 059 log ç 3 + ç Fe è Ferrous ion, molar
1 10–6 10–6
ö ÷ ÷ ø
Ferric ion, molar
E, V-SHE
1 10–12 10–3
+0.771 +0.417 +0.948
58
Corrosion: Understanding the Basics
ferrous ion. The standard electrode potential for this reaction is determined for ferrous ions at 1 M concentration and ferric ions at 1 M concentration. The standard electrode potential for the ferric ion reduction is +0.771 V-SHE. For a ferrous ion concentration of 10–6 M and a ferric ion concentration of 10–12 M, the electrode potential is +0.417 V-SHE. For a ferrous ion concentration of 10–6 M and a ferric ion concentration of 10–3 M, the potential is +0.948 V-SHE. It is apparent that the equilibrium electrode potential for a reaction is quite sensitive to the concentration of the ions involved in the reaction. The equilibrium electrode potential is the value at which the oxidized and reduced species in the reaction are at equilibrium, and the potential at which both can be present and stable at the stated concentrations. This is demonstrated in Fig. 2. For the reaction of ferrous ions being reduced to iron, the standard electrode potential is –0.44 V-SHE. This is for a concentration of ferrous ions equal to 1 M. The significance of this value is that at –0.44 V-SHE, the iron and ferrous ions of 1 M concentration (unit activity) coexist or are at equilibrium. That is to say, with time there will be no net change in the weight of iron immersed in the solution, and there will be no net change in the concentration of ferrous ions in the solution. At a potential more positive than –0.44 V-SHE, the oxidized species or ferrous ion is stable. At a potential more negative than –0.44 V-SHE, the reduced species is stable, that is, iron. The equilibrium electrode potential defines the boundary between the stable oxidized species and the stable reduced species. At a potential value on the boundary, the oxidized species and reduced species are both stable and can coexist. At a potential more positive than the boundary value, the oxidized species is stable and the reduced species reacts to form more of the oxidized species. At a potential more negative than the boundary value, the reduced species is stable and any oxidized species present reacts to form more reduced species.
Fig. 2
Effect of concentrations on the boundary between a stable oxidized species (ferrous ions) and a stable reduced species (iron)
Principles of Aqueous Corrosion
The values calculated for the equilibrium electrode potential at various ferrous ion concentrations are given in Table 4 and are plotted on the electrode potential axis in Fig. 2. As the ferrous ion concentration decreases, the equilibrium electrode potential decreases. The effect of this is that the boundary between the oxidized species (ferrous ion) and the reduced species (iron) moves to more negative electrode potentials. The following behavior is predicted by these calculations. A pure iron rod put into an aqueous environment at –0.71 V-SHE can corrode until the ferrous ion concentration increases to 10–9 M. At this point, the iron and ferrous ion concentration are at equilibrium, and no further net reaction will occur. If the potential of the iron is then made more positive, for example, –0.53 V-SHE, the iron can again corrode and the ferrous ion concentration can increase to 10–3 M. A new equilibrium is established, and the iron is at equilibrium with no net reaction occurring. Now, however, the iron is in contact with a solution whose ferrous ion concentration is equal to 10–3 M. In summary, the calculation of electrode potential based on thermodynamic data and the Nernst equation (Eq 1) is most useful in determination of the boundaries between stable oxidized species and stable reduced species. The equilibrium electrode potential values are a function of the temperature and concentration of oxidized species and reduced species. The Nernst equation provides a sound thermodynamic basis on which the effect of concentration can be determined.
Electromotive Force Series As described in Chapter 2, the electromotive force (emf) series ranks metals with respect to their inherent reactivity. The metals are placed on the series in the descending order of standard electrode potentials. The metals with the most positive values are placed at the top of the series, and the values of electrode potential descend through zero and to larger negative values. The values for standard electrode potential are determined using the Nernst equation for a metal in contact with a solution of its own ions at a 1 M concentration. The values determined under these specific conditions are used to generate the electromotive force series. The standard electrode potential for hydrogen ion reduction to hydrogen gas is defined as the 0.0 reference state by international convention. All other electrode potentials are calculated with respect to this reference state. The standard electrode potentials for the emf series in Table 5 range from 1.5 V versus V-SHE for the reaction of gold ions to gold to a large negative value of –2.37 V-SHE for the reaction of magnesium ions with magnesium. This wide range of potential values indicates a wide range in inherent tendencies to corrode. The more noble metals are at the more positive values. The more active metals, those prone to corrosion, are at the more negative values.
59
60
Corrosion: Understanding the Basics
Table 5
Standard electrode potential for selected metals at 25 °C (75 °F)
Electrode reaction
Au3+ + 3e– « Au Ag+ + e– « Ag Cu2+ + 2e– « Cu 2H+ + 2e– « H2 Pb2+ + 2e– « Pb Sn2+ + 2e– « Sn Ni2+ + 2e– « Ni Cd2+ + 2e– « Cd Fe2+ + 2e– « Fe Zn2+ + 2e– « Zn Al3+ + 3e– « Al Mg2+ + 2e– « Mg
Standard potential, V-SHE
1.50 0.800 0.337 0.000 (Reference) –0.126 –0.136 –0.250 –0.403 –0.440 –0.763 –1.66 –2.37
A more comprehensive listing of electrode potentials can be found in the table of the electromotive force series in Chapter 2.
For any given metal, the standard electrode potential represents the potential at which the metal is at equilibrium with its own ions at a 1 M concentration. At a potential more positive than this value, the metal ion is stable and the metal would react spontaneously to form more metal ions. At potentials more negative than the standard electrode potential, the metal is stable, and no corrosion is predicted. It is important to note that the emf series is for these special circumstances and would not be used by itself to predict the likelihood of corrosion in a practical case.
Galvanic Series The galvanic series is a ranking of metals from the most noble to the most active in a specific environment. Table 6 presents a galvanic series of several commercial metals and alloys in seawater. As in the emf series, the most noble, or cathodic, metals are listed at the top of the series. The ranking proceeds downward to the most active, or anodic, metals. The general positions of metals on the galvanic series and the emf series are similar. Gold and silver are at the most noble end of the series, and magnesium and zinc are at the most active end. However, there are several important shifts of metals in the galvanic series. The principal reason for a shift in the position of a metal or alloy is that some metals and alloys develop stable protective corrosion product films as they react with an aqueous environment. When such a film forms, an alloy is referred to as being in the passive condition. Stainless steels and titanium are two examples of this class of material. The relative position in the galvanic series varies greatly depending on whether the metal is in the active (corroding) or passive (protected) condition. Note in Table 6 the large shift in position for an 18-8 stainless steel (a widely used corrosion-resistant, austenitic 18-8 alloy) in the active condition to a much more noble position in the passive condition. The galvanic series is quite useful in determining the likelihood of detrimental corrosion effects when dissimilar metals are coupled. The
Principles of Aqueous Corrosion
Table 6
Galvanic series of selected commercial metals and alloys in seawater
Noble or cathodic end of series Platinum Gold Graphite Titanium Silver Chlorimet 3 (62Ni-18Cr-18Mo) Hastelloy C (62Ni-17Cr-15Mo) 18-8Mo stainless steel (passive) 18-8 stainless steel (passive) Chromium stainless steel 11-30% Cr (passive) Inconel (passive) (80Ni-13Cr-7Fe) Nickel (passive) Silver solder Monel (70Ni-30Cu) Cupronickels (60-90Cu, 40-10Ni) Bronzes (Cu-Sn) Copper Brasses (Cu-Zn) Chlorimet 2 (66Ni-32Mo-1Fe) Hastelloy B (60Ni-30Mo-6Fe-1Mn) Inconel (active) Nickel (active) Tin Lead Lead-tin solders 18-8Mo stainless steel (active) 18-8 stainless steel (active) Ni-Resist (high-nickel cast iron) Chromium stainless steel, 13% Cr (active) Cast iron Steel or iron Aluminum alloy 2024 Cadmium Aluminum alloy 1100 Zinc Magnesium and magnesium alloys Active or anodic end of series
farther apart two metals are on the galvanic series, the more likely that detrimental galvanic corrosion will occur. When two metals are in electrical contact, the metal at the more noble position in the series becomes the cathode of an electrochemical cell, and the metal in the more active position (lower in the series) becomes the anode and thus corrodes more rapidly. For example, a galvanic couple between platinum and magnesium is disastrous. The more noble platinum greatly accelerates the corrosion of the very active magnesium. Graphite coupled to steel significantly increases the corrosion rate of steel in seawater. When two metals are closer together in the galvanic series, the detrimental galvanic effect is lessened. Several cautions must be recognized when applying the galvanic series. The galvanic series is a ranking of materials for a specific environment under a specific set of conditions. Application of a particular series to a different set of conditions or to a different environment is dangerous and can lead to erroneous predictions. The most common galvanic series is for metals and alloys in nonflowing seawater. Such a ranking system must never be used for predicting corrosion behavior in,
61
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Corrosion: Understanding the Basics
for example, a strongly acidic environment. More detailed information on the galvanic series, including its inherent limitations, can be found in Chapter 4 (in the description of galvanic corrosion).
Standard Electrode Potentials for Other Reactions The emf series presents standard electrode potentials for metals in equilibrium with their own ions. Many practical cases involve additional reactions that are of interest in an operating corrosion cell (see Table 7). Most notable of these reactions are the electrode potentials for the oxygen reduction reactions. Oxygen reacting with hydrogen ions in solution plus electrons can produce water. The standard electrode potential for this reaction is +1.229 V-SHE. An alternative presentation for the reduction of oxygen in alkaline solutions is the reaction of oxygen plus water plus electrons to produce hydroxyl ions. The standard electrode potential for this reaction is +0.401 V-SHE. These electrode potential values along with those presented in the emf series are useful in determining the maximum cell voltage in a corrosion cell made up of two different reactions. The more positive member of the cell is the cathode, and the reaction proceeds in the direction that consumes electrons. The more negative member of the cell is the anode, and the reaction proceeds in the oxidation direction, or the direction that produces electrons. Using values of the standard electrode potentials and the Nernst equation (Eq 1) to determine the effect of ionic concentration on electrode potentials, the electrode potentials for the reactions of interest can be determined. Such determination provides a powerful tool for predicting the likelihood of corrosion and calculating the driving force for the corrosion reactions, that is, the potential difference between the anode and the cathode.
Potential-pH Diagrams: General Aspects One of the most important steps in the science of electrochemical corrosion was the development of diagrams showing thermodynamic conditions as a function of electrode potential (E) and concentration of hydrogen ions, that is, pH. These potential-pH diagrams present a map of the regions of stability of a metal and its corrosion products in aqueous environments. The diagrams identify conditions under which (a) the Table 7
Additional standard electrode potentials
Electrode reaction
O2 + 4H+ + 4e– « 2H2O Fe3+ + e– « Fe2+ O2 + 2H2O + 4e– « 4OH– Cu2+ + e– « Cu+
Standard potential at 25 °C (75 °F), V-SHE
+1.229 +0.771 +0.401 +0.153
Principles of Aqueous Corrosion
63
metal is stable and will not corrode, (b) soluble reaction products are formed and corrosion will occur, and (c) insoluble reaction products are formed and passivity will occur. The diagrams are generated from thermodynamic calculations; they are not determined experimentally by conducting corrosion tests. Potential-pH diagrams are calculated primarily for pure metals in pure water at 25 °C (75 °F). Advances and extensions of the diagrams have been made to include other ionic species and temperatures. These diagrams are also called Pourbaix diagrams in honor of Marcel Pourbaix, a Belgian scientist credited with the development and wide use of these diagrams. Before a discussion of the behavior of specific metals is presented, the format of a potential-pH diagram is described. Figure 3 shows the potential-pH diagram for water with no metal involved. The horizontal axis identifies the pH of the solution. The solutions range from acid conditions at the left (pH 0 to 7) to alkaline solutions at the right (pH 7 to 14). In the middle of the diagram are relatively neutral solutions (pH 5 to 9). The vertical axis is used to plot the oxidizing power (electrode potential) of the solution. These values range from strongly reducing solutions with large negative potentials, for example, –1.2 V-SHE, to strongly oxidizing solutions with large positive potentials, for example, +1.2 V-SHE. The two diagonal lines, identified as (a) and (b), define the region of stability of water as a function of potential and pH. For any value of potential and pH above line (b), water is thermodynamically unstable with respect to the evolution of oxygen. At any conditions of potential and pH below line (a), water is thermodynamically unstable with respect to the generation of hydrogen gas. For potential and pH conditions between lines (a) and (b), water is thermodynamically stable.
Potential, E, V-SHE
(+)
+1.2
(b) O2 + 4H+ + 4e– → H2O 0
(a) 2H+ + 2e– → H2 –1.2
2
Fig. 3
7 pH
12
Potential-pH diagram for water
64
Corrosion: Understanding the Basics
Line (a) represents the equilibrium for the reaction of hydrogen ions to evolve hydrogen gas: 2H+ + 2e– ® H2 Using the Nernst equation, the equation for this line is as follows: E = 0.00 – 0.059 pH The potential at pH 0 is 0.0, and the potential decreases by 0.059 V for each unit increase of pH. At any potential and pH below this line, the hydrogen ion in water will react with electrons to evolve hydrogen gas. Line (b) represents the equilibrium of oxygen plus hydrogen ions and electrons to form water: O2 + 4H+ + 4e– ® H2O Based on the Nernst equation, the equation for this equilibrium is as follows: E = 1.229 – 0.059 pH The potential is +1.229 V-SHE at pH 0 and decreases by 0.059 V for every unit increase in pH. Above line (b) the oxidized species are stable, and water under those conditions reacts spontaneously to produce oxygen and hydrogen ions. The diagram shown in Fig. 3 describes the thermodynamic stability of water and predicts what will occur on a metal surface in contact with water at the full range of potentials and pHs covered by the diagram. If a metal electrode is at a potential and pH below line (a), the water is in an unstable region, and with time, bubbles will evolve on the metal surface. Analysis of these bubbles would show that the gas evolved on the electrode under such conditions is hydrogen. Similarly, for any metal electrode potential-pH condition above line (b), gas evolution will again be observed. Analysis of the gas bubbles under these conditions would show that the evolved gas is oxygen. For any potential-pH condition between lines (a) and (b), water is thermodynamically stable, and no gas evolution would be observed. In dealing with the stability of metals in aqueous solutions, the thermodynamic stability of water is central to discussion of corrosion. Because of this, and to give points of reference, the lines (a) and (b) are included (although not always labeled) on most of the potential-pH diagrams shown in the remainder of this Chapter. The potential-pH diagram can be divided into four regions that describe the general categories of aqueous solutions. As shown in Fig. 4,
Principles of Aqueous Corrosion
65
the diagram can be divided vertically at pH 7. Solutions to the left are acidic, and those to the right are alkaline. The diagram can be further divided along a diagonal line. The upper region is oxidizing, the lower region is reducing. Based on this classification, solutions with potentials and pHs in the upper left of the diagram are characterized as being oxidizing and acidic, and those in the upper right are characterized as oxidizing and alkaline. Those to the lower left are reducing and acidic, and those to the lower right are reducing and alkaline. It is important to note that these are relative descriptions used to compare aqueous solutions with one another. Even though the thermodynamic calculations are based on pure water, potential-pH diagrams can be related to many practical situations with more complex solutions. Figure 5 presents typical pH values for some common solutions, ranging from highly acidic environments, such as battery acid and lemon juice, through mildly acidic solutions, such as boric acid. The figure also shows the neutral solutions, such as milk, through mildly alkaline solutions, such as sodium bicarbonate and borax, and the more highly alkaline solutions, such as sodium hydroxide. The approximate locations of several common environments on the potential-pH diagrams are shown in Fig. 6. Nitric acid (HNO3) is a strongly oxidizing acid, and solutions of HNO3 have potential-pH values (+)
Potential, E, V-SHE
+1.2
0
–1.2
2
Fig. 4 diagram.
7 pH
12
General categories of aqueous solutions. The diagonal line between lines (a) and (b) separates the oxidizing and reducing regions of the
66
Corrosion: Understanding the Basics
pH
Highly alkaline
14
Solution
Sodium hydroxide
Trisodium phosphate
Ammonia
10
Borax Sodium bicarbonate
Milk Neutral
7
Boric acid 5
Tomato
Pickle
Lemon juice
Highly acidic
Fig. 5
1
Battery acid
Range of pH values for some common solutions
in the upper left portion of the diagram. Hydrochloric acid is a strongly reducing acid, and its solutions have potential-pH values at the lower left region of the diagram. Oxygen dissolved in HCl increases the oxidizing power and shifts the potential-pH to somewhat more positive values.
67
Potential, E, V-SHE
Principles of Aqueous Corrosion
pH
Fig. 6
Approximate locations of several common environments on a potentialpH diagram
Sulfuric acid (H2SO4) is a strong acid, yielding low pHs. The oxidizing power of H2SO4 solutions range from mildly reducing to moderately oxidizing. Boric and citric acids represent weak acids and are located to the right of the stronger acids on the diagram. Seawater, fresh water, and atmospheric exposures result in nominally neutral conditions and moderately oxidizing conditions. Bicarbonate and ammonium hydroxide solutions (NH 4 OH) are mildly alkaline and are located just to the right of center in the diagram. Sodium hydroxide (NaOH) forms strongly alkaline environments, and these solutions are located to the extreme right of the diagram. The general behavior of metals in a variety of solutions can often be more easily understood when related to potential-pH diagrams on this basis. The environment determines the general location on the potentialpH diagram. The unique diagram predicts the behavior of a metal in that potential-pH region.
Potential-pH Diagrams for Specific Metals The following paragraphs describe the corrosion (potential-pH) behavior of iron, gold, copper, zinc, aluminum, and titanium. Potential-pH diagrams for some 90 pure elements are available in the Atlas of Electrochemical Equilibria in Aqueous Solution by M. Pourbaix, published by
68
Corrosion: Understanding the Basics
NACE International (see the Selected References at the conclusion of this Chapter). The potential-pH diagram for iron is shown in Fig. 7. The diagram is a map of the regions of immune, passive, and corrosion behavior for iron as a function of potential and pH. Four distinct regions are shown. The region at the bottom of the diagram represents the conditions where iron is immune and no corrosion occurs. Under these conditions, iron is thermodynamically stable. This region covers reducing conditions (negative potentials) across the entire pH range, from acid to alkaline. For any potential-pH combination in this region, iron is stable and does not corrode. The two narrow-banded, cross-hatched areas represent regions where iron corrodes. In both the region to the left (oxidizing and acidic) and the small region to the extreme right (reducing and highly alkaline), iron reacts to form soluble products, and corrosion continues. The shaded area represents a region of passivity for iron. Under oxidizing conditions in neutral to alkaline solutions, iron reacts to form insoluble products and further corrosion reaction is blocked by a protective film. An assumption of the potential- pH diagram is that any insoluble product will be protective. In practice, this is not always the case, and caution must be exercised when using these diagrams for practical applications. For example, the insoluble product on titanium yields a stable passive film. On iron, however, the insoluble products are often porous and not protective in many environments. Potential-pH diagrams allow presentation of a large amount of information in a compact and efficient format. For instance, a single diagram (Fig. 7) describes the general corrosion behavior of iron. Iron is a relatively active metal and corrodes under reducing, moderately oxidizing, and strongly oxidizing conditions in strong acids. In weak acids, iron 2.0
Corrosion
Passivity
0
−1.2
Immunity
2
Potential, E, V-SHE
Potential, E, V-SHE
1.6 +1.2
Fe(OH)3
0.4 Fe2+
0.0
–
HFeO2
−0.4
−1.6 −2
12
Fe(OH)2
Fe
−1.2 0
2
4
6
8
10 12 14 16
pH
pH (a)
Fe3+
0.8
−0.8
Corrosion Corrosion
7
1.2
(b)
Fig. 7
Simplified potential-pH diagrams for iron at 25 °C (75 °F) showing (a) areas of immunity (no corrosion), passivity, and corrosion, and (b) reaction/corrosion products produced
Principles of Aqueous Corrosion
69
Fig. 8
Potential, E, V-SHE
Potential, E, V-SHE
again corrodes under reducing and moderately oxidizing conditions; however, iron can become passive if the oxidizing conditions are increased to a high enough level. Under neutral and mildly alkaline conditions, iron does not corrode, because it either is immune under strongly reducing conditions or is in a passive state for more oxidizing conditions. In strong alkaline environments, iron is free from corrosion except for the small region of potentials and pHs where a soluble, alkaline, corrosion product forms. This general behavior relates fairly well to what is observed in several common environments (Fig. 6). Iron is attacked by strong acids, such as HCl and HNO3. In neutral and mildly alkaline environments, iron is either immune from attack or is protected from corrosion by the formation of passive films. In highly concentrated NaOH solutions, iron exhibits severe corrosion, which is again in agreement with the small region of corrosion predicted by the potential-pH diagram for iron. The specific reactions represented by the lines on the potential-pH diagram for iron are discussed in more detail in the section on construction of these diagrams. The thermodynamically stable species in the immune region is metallic iron. The thermodynamically stable species in the acid corrosion regions are ferrous ions (Fe2+) and ferric ions (Fe3+). Both ferric and ferrous ions are soluble. The thermodynamically stable species for the alkaline corrosion region is hypoferrite ion (HFeO2–). In the passive range, the insoluble corrosion products that are thermodynamically stable are oxides or hydroxides of iron, for example Fe3O4 and Fe2O3. E-pH of Gold. The potential-pH diagrams for gold and iron are shown in Fig. 8. The diagram of iron is repeated in each of these comparisons as a frame of reference showing the relative locations of the regions of stability. It is readily apparent that gold is a much more noble metal than iron. The evidence for this is the wide region of immunity of gold,
(b)
(a)
Potential-pH diagrams for iron and gold. The broad-banded, cross hatched area in the iron E-pH diagram represents a region of passivity. The narrow-banded cross-hatched areas represent where iron and gold will corrode.
70
Corrosion: Understanding the Basics
Fig. 9
Potential, E, V-SHE
Potential, E, V-SHE
extending to the most oxidizing conditions and beyond line (b) on the diagram, which represents oxygen evolution. Gold is thermodynamically stable and will not corrode under a wide range of conditions, from highly acid to highly alkaline and from reducing to strongly oxidizing. The corrosion region for gold extends from highly acid to highly alkaline only at extremely high oxidizing conditions. This is consistent with the observed behavior of gold. Gold is a noble metal and immune from attack, Gold occurs naturally in the surface of the earth in the metallic state. To corrode or dissolve gold requires use of highly oxidizing solutions, such as aqua regia, which is a mixture of concentrated HNO3 (providing a strongly oxidizing environment), and H2SO4 and HCl (to break down surface films.) Iron will corrode readily under a wide range of potential-pH conditions (predominantly acid), while gold is generally immune to corrosion. E-pH of Copper. The potential-pH diagrams for iron and copper are shown in Fig. 9. Copper is a moderately noble metal, with its region of immunity extending above line (a), which represents the line for evolution of hydrogen gas. Copper will corrode under mildly acidic and strongly acidic conditions when the oxidizing power of the environment is moderate to high. Copper will also corrode in strongly alkaline conditions under moderately to highly oxidizing potentials. Copper is resistant to corrosion under neutral to mildly alkaline conditions, because it is immune under reducing conditions and develops protective films (and thus becomes passive) under more oxidizing conditions. Copper is more corrosion resistant than iron under reducing and moderately oxidizing conditions. In a strong acid, iron will corrode in either the deaerated or aerated condition. In contrast, copper is immune from attack under highly reducing, strongly acidic conditions but will corrode under moderately oxidizing, strongly acidic conditions. These conditions are indicated by the arrows on each diagram in Fig. 9. This is consistent with observed behavior. Copper will not corrode in strong HCl acid in the absence of dissolved oxygen or other oxidizing species but will corrode readily in the presence of dissolved oxygen. The addi-
(b)
(a)
Potential-pH diagrams for iron and copper. A, aerated; D, deaerated
71
Potential, E, V-SHE
Potential, E, V-SHE
Principles of Aqueous Corrosion
Fig. 10
(b)
(a)
Potential-pH diagrams for iron and zinc
+
Fig. 11
Potential, E, V-SHE
Potential, E, V-SHE
tion of oxygen to the corrosive environment is sufficient to increase the oxidizing condition and to shift copper from a potential-pH in the immune region to a potential-pH in a corrosion region. E-pH of Zinc. The potential-pH diagrams for iron and zinc are shown in Fig. 10. Zinc is a more active metal than iron. This is indicated by the immune region for zinc being displaced to significantly lower levels than that for iron. Zinc will exhibit active corrosion under acidic conditions ranging from reducing through highly oxidizing and in strongly acidic up through neutral environments. Zinc also sustains attack and forms soluble corrosion products under strongly alkaline conditions. Zinc resists corrosion in mildly alkaline conditions due to the formation of passive films. This is consistent with the observed corrosion behavior of zinc. Zinc, an active metal, is used extensively for sacrificial cathodic protection of iron and for exposed surfaces in atmospheric corrosion, where it develops a passive film. Zinc is not used for structural purposes under strongly acidic or strongly alkaline conditions. E-pH of Aluminum. The potential-pH diagrams for iron and aluminum are shown in Fig. 11. Aluminum is more active than iron or zinc; its immune region is at the lower extremes of the diagram. Aluminum corrodes
+
Potential-pH diagrams for iron and aluminum
(b)
(a)
72
Corrosion: Understanding the Basics
+
Fig. 12
Potential, E, V-SHE
Potential, E, V-SHE
in mildly and strongly acidic and in mildly and strongly alkaline environments. The insoluble corrosion products are stable under neutral conditions and a wide range of oxidizing conditions. This passive range over neutral conditions is responsible for the wide use of aluminum and aluminum alloys for structural purposes. When these alloys are being used, strongly alkaline or strongly acidic conditions must be avoided. E-pH of Titanium. The potential-pH diagrams for iron and titanium are shown in Fig. 12. Titanium is an extremely active metal, and its immune region is nearly off the bottom of the potential-pH diagram. The excellent corrosion resistance of titanium and titanium alloys in a wide range of solutions can be attributed to the tenacious, passive film that forms on them. The passive region extends over the entire range of pH, from highly acidic to highly alkaline, under moderately reducing to highly oxidizing conditions. Titanium does have a corrosion region under reducing and highly acidic conditions, as shown at the lower left region of the diagram. Titanium will corrode under these conditions, as demonstrated by the attack it undergoes in warm deaerated hydrochloric acid. This environment is highly acidic and strongly reducing and results in the breakdown of the protective film that forms on titanium, followed by significant corrosion. Comparison of E-pH Diagrams for Metals. The potential-pH diagrams for several metals in water at 25 °C (75 °F) are shown in Fig. 13. Metals exhibit a wide range of corrosion behaviors, extending from the noble alloys—such as gold, iridium, and platinum—to extremely active metals—for example, beryllium, zinc, and aluminum. Some of the active metals can be used in many corrosion applications because of the formation of insoluble corrosion products and the development of protective passive films. It is important to realize, however, that it is the integrity of these protective films that is being relied on for the corrosion resistance of these materials. A much more reliable practice is the application of immune metals, where corrosion resistance is imparted by thermodynamic stability.
+
Potential-pH diagrams for iron and titanium
(b)
(a)
Principles of Aqueous Corrosion
73
The behavior of several of these metals in specific environments can be compared by first identifying a specific environment, labeled A to H in Table 8. Environment A represents a strongly acidic (pH 0) and strongly reducing (0 V-SHE) condition. The location of environment A on the potential-pH axis is indicated on the diagram in the upper right of Table 8, along with locations of the other environments. The corrosion behavior of various metals can be determined in each of the environments by comparing the environment location with the predictions of the potential-pH diagrams presented in Fig. 13. Taking aluminum as an example, corrosion is predicted in environments A, B, and C, and passivity is predicted in environments D and E. Corrosion is also predicted in environments F, G, and H. Similar comparisons and tabulations can be made for titanium, copper, iron, zinc, and gold. The comparison of the corrosion behavior of a 0
7
14
0
7
14
0
7
14
0
7
14 +2
+1
+1
0
0
E, V-SHE
+2
–1
–1
–2 0
Gold
E, V-SHE
0
7
14
0
14
0
Platinum 7
14
0
14
0
Copper 7
14
0
14
0
Cobalt 7
–2 14 pH 14
+2
+2
+1
+1
0
0
–1
–1
–2 0
E, V-SHE
0
Cadmium 7
14
0
14
0
Iron 7
14
0
14
0
Zinc 7
14
0
Aluminum
–2 14 pH
14
+2
+2
+1
+1
0
0
(a) (b) (c) (d)
–1
–1 –2 0
Zirconium
Fig. 13
14
0
Titanium
14
0
Beryllium
(e)
–2 14 pH
Potential-pH diagrams for several metals in water at 25 °C (75 °F). (a) Corrosion by dissolution. (b) Corrosion by gasification. (c) Passivation by oxide or hydroxide layer. (d) Passivation by hydroxide layer. (e) Immunity
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Corrosion: Understanding the Basics
Table 8 Corrosion behavior of metals in a variety of environments based on potential-pH diagrams Environment
pH, conditions
A ® strong acid B ® strong acid C ® strong acid D ® neutral E ® neutral F ® strong alkaline G ® strong alkaline H ® mild alkaline
Location on E-pH
E, V-SHE
0, strongly reducing 0, mildly oxidizing 0, strongly oxidizing 7, reducing 7, oxidizing 14, strongly reducing 14, strongly oxidizing 10, mildly oxidizing
Oxidizing
0 0.5 1.0 –0.3 0.5 –0.8 –0.2 0
+2
C B A
+1
E
0
E D
G
H
F
–1 Reducing
–2 0
7
14
pH– Environment
pH
A B C D E F G H
0 0 0 7 7 14 14 10
E, V-SHE
Al
Ti
0 +0.5 +1.0 –0.3 +0.5 –0.8 +0.2 0
C C C P P C C C
P P P P P P P P
Corrosion behavior(a) Cu Fe
I C C I P I C P
C C C C P C P P
Zn
Au
C C C C C C C P
I I I I I I I I
(a) C, corrosive; P, passive; I, immune
material in any given environment can be readily determined by reading across Table 8 and identifying regions of corrosion, passivity, or immunity.
Strategies for Corrosion Control from E-pH Diagrams The potential-pH diagram for iron in water at 25 °C (75 °F) can be used to rationalize methods of corrosion control. The starting position for the analysis will be iron in a mildly oxidizing acid condition; that is, E = –0.2 V-SHE and pH = 4. This potential-pH location is indicated by the cross in Fig. 14, a condition representative of iron in a mildly oxidizing, weak acid. Under this condition, iron is predicted to corrode with formation of soluble ferrous ions. Corrosion can be controlled by employing several options, indicated by the arrows and letters in Fig. 14. If the pH of the environment is increased to neutralize the acid and to move into the slightly alkaline region, the iron is moved from a region of corrosion by the formation of soluble products to a region of passivity through the formation of insoluble products. This method of corrosion control is a form of water treatment or change in environment and is indicated by path A. The corrosion of iron can also be controlled by making the conditions more reducing. This is done by changing the potential of the steel surface from mildly oxidizing to a more negative potential. At the more reducing potential, iron is in an immune region and is thermodynamically stable. This control method is cathodic protection and involves the application of a protective current that shifts the potential of the steel surface to more negative values, as indicated by path B in Fig. 14.
Principles of Aqueous Corrosion
Fig. 14
Methods of corrosion control for iron related to the potential-pH diagram
Another method for controlling the corrosion of iron under this environment would be to increase the oxidizing conditions. This method of protection, that is, anodic protection, requires the movement of the steel electrode potential from a region of corrosion resulting from soluble ion formation into a region of more oxidizing conditions where passivity results from the formation of insoluble corrosion products. This method of control is represented by path C in Fig. 14. It is important to understand the difference between anodic protection (path C) and cathodic protection (path B). In cathodic protection, the steel is moved into a region of immunity where iron is the thermodynamically stable phase. In anodic protection, the iron is moved into a region where passivity is favored. If the protective nature of the insoluble products is distributed such that they do not provide complete coverage or if the anodic protection is incomplete, then corrosion can in fact be made much more severe. Both of these methods of corrosion protection are described in Chapter 10. The corrosion protection methods indicated by methods D and E in Fig. 14 involve increasing and extending the potential-pH range over which passive behavior is observed. This can be accomplished by the addition of a passivating chemical species, such as chromate ions, to the environment, promoting the formation of insoluble products on the surface (path D). Alternatively, the region of passivity can be extended by adding constituents to the iron that promote the formation of insoluble films. This alloying option for iron includes the addition of chromium (Cr), nickel (Ni), and molybdenum (Mo). Iron alloyed with more than 12% Cr, plus Ni and Mo, has spontaneous passivity over a wide range of environments. These alloys constitute entire families of austenitic stainless steels and nickel alloys. This general approach to the application of potential-pH diagrams for determining and identifying corrosion control strategies can be quite
75
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Corrosion: Understanding the Basics
useful. Each control method can be evaluated based on the initial potential-pH condition for the metal and environment.
Limitations of E-pH Diagrams The preceding discussion emphasizes the utility and wide application of potential-pH diagrams. However, their limitations must also be understood. Potential-pH diagrams are the result of thermodynamic calculations and do not provide any kinetic data or information as to reaction rates. The thermodynamically stable state and species are identified based on the lowest free-energy state for the reactions considered. There are two dangers. The first danger that the thermodynamically stable species will not be achieved due to kinetic considerations. The second danger is that operative reactions might not be considered due to oversight or the lack of complete knowledge of the constituents in a complex environment. The potential-pH diagrams are based on the reaction of pure metal with the constituents of water. This approach neglects all the possibly important effects of other ions in solution, for example, chloride, nitrate, reduced sulfur species, and beneficial species such as chromate ions. The diagrams presented in this Chapter are calculated for 25 °C (75 °F). Changes in temperature can greatly affect the regions of passivity and shift the regions of immunity. The treatment of potential-pH diagrams has dealt specifically with calculations for pure metal, while most engineering applications deal with alloys or mixtures of metals. The treatment assumes that all insoluble products result in passivity. In practice, this is not the case. Many insoluble products are porous or loosely attached to the metal surface and provide limited protection against further corrosion. For example, oxides on titanium are quite protective, whereas oxides on iron can be either protective or nonprotective, depending on the specific conditions. Yet another limitation is that use of potential-pH diagrams generally involves the prediction of behavior at a given potential and pH. As the corrosion reaction proceeds, it is likely that the potential and pH will change from their initial values. The trajectory of potential-pH variance is not readily determined without employing additional information. Although these limitations are substantial, the use and benefit of potential-pH diagrams in the understanding and control of corrosion cannot be disputed. Extensions of the treatment presented in this Chapter have provided for considerable progress in overcoming several of the limitations previously noted. One of the drawbacks to many of these extensions, however, is that the major benefit of the potential-pH diagrams is often sacrificed. That is, the ultimate simplicity and convenience of presentation of a large amount of information in a compact
Principles of Aqueous Corrosion
77
and efficient form are compromised when other considerations are added to the analysis.
The Kinetics of Aqueous Corrosion As described in the first half of this Chapter, thermodynamics has been widely applied to corrosion studies for many years. Use of thermodynamic calculations permits prediction of conditions under which a metal is stable and corrosion will not occur. Thermodynamics does not predict, however, how fast the corrosion reaction will occur. To understand the kinetics or rate of corrosion, examination of a number of electrochemical principles for example, mixed-potential behavior, polarization behavior, and the concept of exchange currents, for example, is required.
Electrochemical Reactions The characteristics of an electrochemical reaction are that oxidation (or the generation of electrons) and reduction (or consumption of electrons) both occur, and there is an electron transfer from the anode to the cathode. The characteristics of an electrochemical reaction are illustrated by considering the behavior of iron in hydrochloric acid. Iron reacts vigorously with HCl; hydrogen is evolved and the iron gradually goes completely into solution. The reaction is: Fe + 2HCl ® FeCl2 + H2
(Eq 4)
The solid iron gradually disappears and a gas is evolved. This can be seen with the naked eye. The solution in the above reaction is ionized and contains positively and negatively charged ions. The HCl contains hydrogen ions (H+) and chloride ions (Cl–). One drop of acid contains millions of each of these ions. Likewise, ferrous chloride in solution can be considered as iron ions (Fe2+) and chloride ions (Cl–). Therefore, Eq 4 can be written as the following: Fe + 2H+ + 2Cl– ® Fe2+ + 2Cl– + H2
(Eq 5)
The iron converted to an iron with two positive charges. By definition, the iron is said to have been oxidized (loss of electrons). On the other hand, the hydrogen ions have each gained an electron. By definition, they have been reduced. The overall reaction can be considered as two separate ones: Oxidation: Fe ® Fe2+ + 2e–
(Eq 6)
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Corrosion: Understanding the Basics
Reduction: 2H+ + 2e– ® H2
(Eq 7)
Equations 6 and 7 both take place on the surface of the metal. The areas where oxidation occurs are defined as anodes, and those where reduction takes place are defined as cathodes. An electrical potential exists between the anode and cathode areas. The electrons produced in Eq 6 flow through the metal to the cathode areas to take part in the reaction of Eq 7. Hydrogen ions in the vicinity of the anode areas are not needed there, and they flow (under the influence of the potential difference) to the cathode to sustain the reduction reaction. A complete electrical circuit exists, and a current flows from anode to cathode. The faster the solid is converted to iron ions (i.e., the greater the corrosion), the larger is the current flowing in this corrosion cell. The mechanism for the preceding case is shown schematically in Fig. 15. Many such corrosion cells occur on a corroding metal surface. Anode and cathode sites can switch roles so that uniform corrosion can occur. Equations 6 and 7 must occur at the same rate to conserve the electrical neutrality of the metal. If Eq 6 stops, then so must Eq 7, and vice versa. The corrosion reactions of other metals and corrodents are similar. In all cases of electrochemical corrosion, the anodic reaction consists of oxidation of the solid metal, M, to its ions in solution as follows: M ® Mn+ + ne– When an alloy corrodes, several anodic reactions can take place simultaneously. For example, the anodic reactions for an aluminum-copper alloy might be the following: Al ® Al3+ + 3e– Cu ® Cu2+ + 2e– Mg ® Mg2+ + 2e–
H
+
Hydrochloric acid (H+ and Cl–) H2
Fe2+ + 2e– 2H+ + 2e–
Anode
e Fe
Fig. 15
Cathode
–
Corrosion of iron in hydrochloric acid
Solid iron (Fe)
Principles of Aqueous Corrosion
79
Therefore, the possible anodic reactions in a system are relatively easy to predict. With cathodic reactions, there are more possibilities. Various types of reduction can occur, including the following, for example: Hydrogen evolution in acid solutions: 2H+ + 2e– ® H2 Oxygen reduction in acid solutions: O2 + 4H+ + 4e– ® 2H2O Oxygen reduction in neutral or alkaline solutions: O2 + 2H2O + 4e– ® 4OH– Metal-ion reduction: M3+ + e– ® M2+ Metal plating: M+ + e– ® M Hydrogen evolution and oxygen reduction are the most common cathodic reactions. Different cathodic reactions can occur simultaneously on a corroding metal surface. In summary, corrosion in the electrochemical sense occurs by solid metal being oxidized to positively charged metal ions in solution. This occurs at areas called anodes. The resulting excess of electrons passes through the metal to surface areas called cathodes where electrons are removed by a reduction reaction. The corrodent must contain a species that can be reduced at the cathode and ions capable of completing the electrical circuit between anode and cathode areas (i.e., the corrodent must be an electrolyte). A current flows through the solution from anode to cathode areas. As described subsequently, the driving force in the reaction is the electrical-potential difference that causes current flow between the anode and cathode.
Mixed-Potential Theory Mixed-potential theory provides a solid basis for the treatment of the rates of electrochemical reactions. The concepts of mixed-potential theory can be used to describe and explain many of the material and environmental effects of observed in corrosion phenomena. The theory
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Corrosion: Understanding the Basics
particularly delineates the importance of the reduction reactions in an electrochemical process as well as the anodic and dissolution process. Two basic premises underlie mixed-potential theory. First, the anodic current in an electrochemical cell must equal the cathodic current. This is a requirement of the conservation of electrical charge. The number of electrons generated by the total of all oxidation reactions must be exactly equal to the number of electrons consumed by the total of all reduction reactions. Second, for the purpose of examination and definition of the electrochemical cell, the anodic (oxidation) and cathodic (reduction) reactions can be defined, written, and treated independently. As shown in Fig. 15, the anodic reactions account for the generation of electrons, and the cathodic reactions account for the consumption of electrons. Conceptually, these can be dealt with separately. For an operating corrosion cell, then, the reactions are coupled through the principle of conservation of electrical charge. The application of mixed-potential theory to corrosion phenomena is discussed in succeeding paragraphs. More detail is available in several of the selected references listed at the conclusion of this Chapter. The free-energy change for the oxidation of iron to ferrous ions is shown schematically in Fig. 16. As the free-energy wells are drawn in Fig. 16, the system can lower its free energy, DG, by the oxidation of an iron atom at the metal surface to form a dissolved metal ion in solution. The net decrease in free energy is indicated by DG oxidation. The free energy of the reaction is related to the electrochemical potential through the Nernst equation (Eq 1). The net free-energy change, DG oxidation, is equivalent to an electrochemical overpotential (hanodic). The reaction in Fig. 16 shows a spontaneous reaction of iron to corrode (oxidize) with a net decrease in free energy or a positive overpotential.
Fig. 16
Free energy and electrochemical potential for an activation-controlled reaction. Reaction rate increases exponentially with driving force.
Principles of Aqueous Corrosion
81
The driving force for this corrosion reaction is the free-energy change or the overpotential. The rate of the reaction can be expressed as a current or the number of electrons generated per unit time. The rate of the reaction (current) increases exponentially with the driving force (overpotential): æ –DG ö Current = A expç ÷ è RT ø æ nFh ö Current = A expç ÷ è RT ø where A is a constant, DG is the free-energy change, R is the gas constant, T is temperature, n is the number of electrons generated per ion, F is Faraday’s constant, and h is overpotential (DE). The overpotential is equivalent to DE or E–E0, where E is the potential of the metal surface and E0 is the equilibrium potential for the reaction. An equivalent expression for this relationship is that the overpotential is proportional to the log of the current, that is, h µ log i. This relationship indicates that a unit increase in the driving force, that is, the overpotential, results in an order of magnitude increase in the reaction rate, that is, current. An electrochemical reaction that behaves as described in the preceding paragraph is referred to as being under the control of activation polarization. Activation polarization simply means that the driving force for the reaction (overpotential) is proportional to the log of the reaction rate (current). The relationships for anodic and cathodic processes under activation polarization are as follows: ha = a a + b a log i (anodic ) hc = a c + b c log i ( cathodic ) Anodic overpotential (ha) equals a constant aa plus a constant ba times the logarithm of the current. The cathodic overpotential (h c) is equal to a constant ac minus a constant bc times the logarithm of the cathodic current. The constants a and b are referred to as the anodic and cathodic Tafel constants, respectively. Over the potential range where these equations describe the relationships between potential and current, reaction is under activation control. Mixed-potential theory deals with the anodic and the cathodic reactions independently. The reactions then are coupled by the rule that the total of all anodic currents (ia) and the total of all cathodic currents (ic) are equal. Where activation control governs the reaction rate, there is an exponential relationship between potential and current.
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Corrosion: Understanding the Basics
Types of Polarization As stated in the preceding sections, the driving force for the corrosion reaction is the potential difference between the anode and the cathode in the corrosion cell, and the reaction rate is equal to the current that flows through the cell. The total resistance throughout the corrosion cell is the total of the resistance associated with the anode plus the resistance of the ionic path plus the resistance associated with the cathode plus the resistance of the electronic path. As with Ohm’s law, that is, E = iR, for a simple electrical circuit, the potential, current, and resistance in a corrosion cell are related. For a given potential difference between the anode and the cathode, the current or corrosion rate will increase as the resistance throughout the cell decreases or, conversely, the corrosion rate will decrease as the resistance throughout the cell increases. Various types of polarization describe the resistance elements throughout the corrosion cell. Three forms of polarization are discussed in this section: activation polarization, concentration polarization, and ohmic polarization. Activation polarization is defined by the relationships given in the previous section describing electrochemical reactions. Figure 17 graphically illustrates the relationship for the anodic reaction, which is the oxidation of iron atoms to soluble, ferrous ions. The potential is plotted on the vertical axis, with more positive or oxidizing potentials at the top. The reaction rate or log of the anodic current is plotted on the horizontal axis, with increasing currents to the right. The activation polarization curve for this anodic reaction is a straight line extending from the lower left to the upper right. The slope of the line is given by the anodic Tafel slope, ba. Figure 17 shows that a unit increase in the overpotential, E, results in an order of magnitude increase in the reaction rate. As the oxidizing potential becomes more positive, the rate of the corrosion or oxidation of iron increases, that is, the anodic overpotential increases.
Fig. 17
Activation polarization curve for the anodic reaction of iron and ferrous ions
Principles of Aqueous Corrosion
83
Ferrous ions are generated more rapidly, and more electrons are left at the metal surface. The activation polarization curve for the cathodic reaction of hydrogen ions to form hydrogen gas is shown in Fig. 18. The axes are the same as described for Fig. 17. The cathodic or reduction reaction is a straight line extending from the upper left to the lower right. The slope of the line is the cathodic Tafel slope, bc. The reduction reaction increases as the oxidizing conditions become more reducing or the overpotential becomes more negative. As the potential becomes more reducing by a unit value, the rate of the reduction reaction increases by an order of magnitude. As the conditions become more reducing, hydrogen ions and electrons are consumed more rapidly at the metal surface and more hydrogen gas is generated. It is important to note that the anodic current, ia, is the rate of generation of electrons and that cathodic current, ic, is the rate of consumption of electrons. In an operating electrochemical cell, oxidation current leaves the metal surface at the anode, and a cathodic current enters the metal surfaces at the cathode. The anodic reaction of Fig. 17 and the cathodic reaction of Fig. 18 are combined on a single diagram in Fig. 19—an activation polarization diagram for the anodic and cathodic reactions. The requirements of mixed-potential theory are met at only a single point, that is, the point where the anodic and cathodic reaction curves cross. This is the only location at which the anodic reaction rate equals the cathodic reaction rate. The potential of this intersection is identified as Ecorr, and the current at this intersection is defined as icorr. The potential intersection is referred to by several terms, including the corrosion potential, the free corroding potential, and the open-circuit potential. The current intersection is referred to as the corrosion current. At potentials more positive or more oxidizing than the corrosion potential, the anodic current is greater than the cathodic current, and more electrons are generated than (+) 2H+ + 2e –
H2
Cathodic reaction:
= c – c log ic
E c
(–) Log ic
Fig. 18
Activation polarization curve for the cathodic reaction of hydrogen ions and hydrogen gas
84
Corrosion: Understanding the Basics
are consumed. At potentials more negative or more reducing than the corrosion potential, the cathodic current is greater than the anodic current, and more electrons are consumed than are generated. A steady state of no net consumption or generation of electrons is achieved only at the corrosion potential. Neither of these cases (i.e., potentials more oxidizing or more reducing than the corrosion potential) results in a steady state. In order to maintain a system away from the corrosion potential (Ecorr), current must be supplied from an external source or other reactions. The slopes ba and bc are determined by the properties of both the metal surface and the electrolyte. There are three forms of activation polarization control: mixed control, cathodic control, and anodic control (Fig. 20). The type of control essentially results from the slope values of the anodic and cathodic curves. Under mixed control, the corrosion rate is equally sensitive to shifts in the anodic or oxidation reaction and the cathodic or reduction reaction. Under cathodic control, the slope of the reduction curve is greater than the slope of the oxidation curve. This re(+)
Anodic
Combined reactions at steady state:
ianodic = icathodic
Ecorr E
Cathodic
icorr
(–) Log i
Fig. 19
Combined diagram of an anodic reaction and a cathodic reaction with activation polarization.
Fig. 20
Schematic diagrams of the three forms of activation polarization control
Principles of Aqueous Corrosion
sults in the corrosion reaction being more sensitive to changes in the reduction reaction kinetics than to changes in the oxidation reaction kinetics. Under anodic control, the slope of the oxidation reaction is greater than the slope of the reduction reaction curve. This results in the corrosion rate being more sensitive to changes in the anodic reaction kinetics than to changes in the cathodic reaction kinetics. It is useful to note that form of control which is operative under a given set of circumstances. This knowledge allows one to focus on control of the most sensitive portion of the corrosion cell; that is, corrosion control can be focused on either the anodic or the cathodic reaction, whichever will have the most effect on lowering the corrosion rate. In practice, systems employing all three types of activation control are used. Concentration polarization results from mass transfer or diffusionlimited effects at the reacting metal surface. The effect of concentration polarization on the shape of the cathodic polarization curve for a reduction reaction is shown in Fig. 21. At lower currents, the cathodic reaction is controlled by activation polarization, and a straightline sloping from the upper left to the lower right results. The slope of this curve is –bc, the cathodic Tafel slope. In this region, as the potential becomes more negative or more reducing, the rate of the reaction increases. As the potential continues to become more reducing, the overpotential of the reduction reaction increases, as does the driving force for the reduction reaction. There is, thus, a greater driving force for reduction, and the rate continues to increase. As the cathodic overpotential becomes even greater (i.e., the potential becomes even more negative), a point is reached at which increasing the driving force no longer increases the rate of reaction. This is the region of concentration polarization.
Fig. 21
Onset of concentration polarization at more reducing potentials for a cathodic reduction reaction
85
86
Corrosion: Understanding the Basics
The dashed horizontal line on Fig. 21 indicates the separation between the potential region where the reaction is under activation control (above the dashed line) and the potential region where the reduction reaction is under concentration control (below the dashed line). The region of activation control is described by the straight, sloped line, and the concentration control region is described by the vertical line. In the region of concentration control, there is more than sufficient driving force for the reaction to occur as reactants arrive at the electrode surface. Therefore, the reaction rate is no longer dependent on the potential (driving force) but rather becomes dependent on either the rate at which reactants species arrive at the electrode or the rate at which products of the reaction can be removed from the vicinity of the electrode surface. The reaction rate is controlled by the mass transfer of reactant and products of the reaction, and this mass transport is essentially independent of potential. A curved portion of the reduction reaction kinetics curve is shown near the horizontal line. In this region, both the activationdependent kinetics and concentration-control kinetics affect the potential versus log current behavior. In the linear regions of the curve, the kinetics are dominated either by activation control or by concentration control. The magnitude of the reduction current in the region of concentration control, that is, the value of current where the vertical segment of the curve intersects the log current axis, is defined as the limiting current (iL), which can be determined by the following:
(
)
iL = DnF c d
where D is the diffusion coefficient for the ionic species being reduced at the electrode surface, n is the number of electrons consumed by the reduction reaction, F is Faraday’s constant, d is the thickness of the diffusion layer of the electrolyte near the electrode surface, and c is the concentration of the species being reduced. This relationship describes the limiting current for the reduction kinetics, where the rate of reaction is controlled by the arrival of the species being reduced at the electrode surface. For example, in the hydrogen evolution reaction this reaction will predict the reaction rate when the arrival of hydrogen ions from the solution to the metal surface controls the reaction rate. The effect of increasing the magnitude of the limiting current is shown by the dashed lines in Fig. 21. As the increasing limiting current moves the vertical segment of the line farther to the right to higher current values, more of the linear portion controlled by activation control is observed. Referring to the equation above, the factors that can result in increasing limiting current are identified. The limiting current under concentration control will increase when the diffusion coefficient of the
Principles of Aqueous Corrosion
species being reduced (D) increases, when the concentration of the species being reduced (c) is increased, and when the thickness of the diffusion layer (d) is decreased. The preceding discussion has focused on the concentration-control portion of the reaction rates for a reduction reaction. Concentration control also can be observed for oxidation or anodic reactions. For such reactions, concentration polarization is delineated where the anodic polarization curve on a diagram of potential versus log current deviates from a straight, diagonal line with a slope equal to the anodic Tafel slope and approaches a vertical line that is independent of potential. Ohmic Polarization. The third form of polarization in an electrochemical cell is ohmic polarization, which results from pure resistance elements along the current path in the cell. Ohmic resistance is also referred to as iR effects. Ohmic polarization is observed either in the region of ionic conductivity, where current is transported by the movement of ions through the electrolyte from the anode to the cathode, or in the electronic conductivity region, where current is transported through the metallic path from the cathode to the anode. Because metals have relatively low resistivity or high conductivity, ohmic polarization along the electronic path is not generally significant. Exceptions to this include cases where the current travels long distances between the cathode and anode along the metallic path. Ohmic polarization effects are not uncommon in the electrolyte, because many electrolytes have significant resistances to the ionic current flow. The effect of ohmic polarization on the corrosion current in an electrochemical corrosion cell is shown in Fig. 22. Three cases are shown, with the resistance of the solution increasing from a value of R1 (where resistance is essentially 0) to R3 (a high resistance value). The resulting current through the corrosion cell for each of these resistances is indicated by i1, i2, and i3. When the conductivity of the solution is quite high, corresponding to essentially no resistance in the electrolyte, the cathodic reduction curve and the anodic oxidation curve intersect at i1. In this case, the potentials of the anode and cathode are polarized to the same value. As the resistance of the solution increases, the potential of the anode and the cathode are no longer equal. A potential drop (iR) in the solution results from the passage of current (i) through the resistive solution (R). The resulting effect is that some of the potential difference between the anode and cathode is taken up by the potential drop through the solution and is, therefore, unavailable to drive the activationcontrolled reactions. As the resistance of the solution increases from R1 to R2 to R3, the magnitude of the potential drop of ohmic polarization in the cell increases, with a resulting decrease in the corrosion current of the cell from i1 to i2 to i3. Note that in all cases the anodic current equals the cathodic current.
87
88
Corrosion: Understanding the Basics
(+)
Ec
iR
E
i1
i2 Ea
i3 Log i
R3 > R2 > R1 = 0
Fig. 22
Effect of ohmic polarization on the current in a corrosion cell
The effect of ohmic polarization on the corrosion rate can be illustrated by the behavior of galvanic couples comprised of copper electrically coupled to steel in different waters. The driving force for the corrosion reaction is the potential difference between the copper cathode and the steel anode. In seawater, which has high conductivity, there is very little resistance to current flow, and a high corrosion rate (equivalent to i1) is observed. Tap water has a much lower conductivity than seawater, and the resulting corrosion rate is also significantly lower (equivalent to i2). Distilled water has even less conductivity than tap water, and the resulting corrosion rate is still lower for the galvanic couple (equivalent to i3). Thus, an effective way to reduce the corrosion current in a corrosion cell is to increase the resistance to ionic current flow through the cell.
Applications of Mixed-Potential Theory Diagrams The basis for mixed-potential theory diagrams and the various forms of polarization that can be observed and described using these diagrams were described in the preceding sections. These diagrams can be most useful for predicting and explaining observed corrosion behavior. Examples of the application of mixed-potential theory are presented here. Corrosion of Zinc in Deaerated Acid. Figure 23 describes the corrosion behavior of zinc, an active metal, in a strong (pH 0) deaerated acid. In this system, two half-cell reactions are coupled in an electrochemical corrosion cell; thus, the principles of the mixed-potential theory will be obeyed. The system will be polarized to a potential where
Principles of Aqueous Corrosion
89
i0
+0.2
E, V(SHE)
0
EH / H+ 2 –0.2
icorr –0.4
Ecorr i0
–0.6
EZn/Zn2+ –0.8 10–12
10–10
10–8
10–6
10–4
10–2
Current density, amp/cm2
Fig. 23
Mixed-potential theory diagram for the corrosion of zinc in a deaerated acid
the sum of the currents from the cathodic reduction reactions equals the sum of the currents from the anodic oxidation reactions. The opencircuit potential and current densities (i0) for each half-cell reaction are shown in Fig. 23. The corrosion potential, Ecorr, and the corrosion current, icorr, for the coupled reactions in the corrosion cell are also indicated. When zinc is immersed in a strong acid, the corrosion potential is approximately –0.5 V-SHE, and the corrosion current is approximately 10–4 A/cm2. The oxidation current at the anodic results almost completely from the oxidation of zinc atoms and the formation of zinc ions in solution, with the release of electrons at the metal surface. The reduction reaction current is almost completely the result of the reduction of hydrogen ions from solution and the formation of hydrogen gas, with the consumption of electrons at the metal surface. At the corrosion potential, zinc ions are being produced and hydrogen ions are being consumed at a rate equivalent to the corrosion current. The mixed-potential theory diagrams for the two half-cell reactions can, thus, be used to predict the resulting corrosion behavior of zinc in a deaerated acid. Unless an external current is impressed upon the system, the system will remain at the steady-state condition, where the sum of the anodic currents equals the sum of the cathodic currents. Effect of pH on Corrosion of Active Metals. The effect of increasing the acidity of the solution on the corrosion rate of an active metal is shown in Fig. 24. As the acidity is increased (lower pH), the hydrogen ion concentration in solution increases where the reduction reaction is due to the reduction of hydrogen ions and evolution of hydrogen gas, increasing the concentration of hydrogen ions in solution has the effect of increasing the cathodic reduction rate at every potential. The effect is shown in Fig. 24 by shifting the reduction curve to the right from Ec(1)
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Corrosion: Understanding the Basics
to Ec(2). As the reduction kinetics are increased, the intersection with the anodic dissolution curve also shifts to higher current values. The corrosion rate for the higher-pH solution (lower hydrogen ion concentration) is i1, and the corrosion rate in the lower-pH solution (higher hydrogen ion concentration) is i2. Therefore, for the corrosion of an active metal in a deaerated acid, the corrosion rate of the metal is predicted to increase as the strength of the acid increases. Active-Passive Behavior. Anodic polarization curves for an active metal and an active-passive metal are shown in Fig. 25 (the basic concepts associated with passive behavior are described in Chapter 2). The anodic dissolution rate of an active metal increases as the potential becomes more oxidizing. A linear increase is observed until the onset of concentration polarization. The behavior of the active-passive metal is similar at the start; that is, as the potential increases, a linear increase is observed. However, as the potential continues to become more oxidizing, a sharp drop to a much lower corrosion current is observed. This sharp decrease in current corresponds to the attainment of a potential range in which an insoluble corrosion product forms, significantly lowering the rate at which the metal corrodes. Within this passive range, the small corrosion current is independent of potential, as indicated by the vertical line at higher potential for the active-passive metal. A polarization curve typical of the active-passive behavior of stainless steels is shown in Fig. 26. The effect of increasing the rate of cathodic reaction on the corrosion behavior of an active-passive metal is shown in Fig. 27. For the
Fig. 24
Effect of lower pH (greater acidity) on the corrosion rate of an active metal
Principles of Aqueous Corrosion
91
initial conditions, the corrosion potential and corrosion current are indicated by Ecorr(1) and icorr(1), respectively. The reduction kinetics curve intersect the oxidation kinetics curve in the active region, and a relatively high corrosion rate is predicted. When the rate of cathodic reaction is increased, as shown by a shift of the cathodic reduction curve to
Fig. 25
Anodic polarization curves for an active metal and an active-passive metal
Potential E, V-SHE
Transpassive
Passive
M
M2++ 2e– Active M+ + e–
M
io(Me/Me2+) Log current density, i
Fig. 26
Schematic polarization curve for a metal (e.g., stainless steel) that displays an active-passive transition. At relatively low potential values, within the active region, the behavior is linear, as it is for normal metals. With increasing potential, the current density suddenly decreases to a very low value, which remains independent of potential. This is termed the passive region. Finally, at even higher potential values, the current density again increases with potential in the transpassive region.
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Corrosion: Understanding the Basics
the right, the intersection is moved to values of Ecorr(2) and icorr(2). The net effect of increasing the rate of cathodic reaction is to significantly decrease the corrosion current in the cell and to shift the corrosion potential to a more positive (more oxidizing) value. Comparison of Fig. 24 and 27 shows that increasing the cathodic reaction rates for an active metal increases the corrosion rate; for an active-passive metal, on the
Fig. 27
Effect of increasing the rate of cathodic reaction on the corrosion behavior of an active-passive metal
Fig. 28
Effect of increasing the concentration of reducible species on the cathodic polarization curve and the corrosion rate of an active metal
Principles of Aqueous Corrosion
other hand, increasing the cathodic reaction rate can significantly decrease the corrosion rate. As with the active metal, for the active-passive metal, the steady-state condition is the point at which the cathodic reaction curve intersects the anodic behavior curve. Effect of Increasing the Concentration of Reducible Species. Figure 28 illustrates the effect of increasing the concentration of reducible species on the cathodic polarization curve and the corrosion rate of an active metal. As the concentration of reducible species increases from c1 to c2 to c3, the cathodic reduction current increases in both the activation-controlled and concentration-polarization-controlled portions of the curve. The net effect is to shift the entire curve to the right. The intersection of the cathodic polarization curve and the anodic polarization curve moves to the right from i1 to i2 to i3 as the concentration of reducible species increases. In each of these cases, the concentration polarization portion of the curve, that is, the limiting current value, intersects the anodic polarization curve. The prediction of the mixed-potential theory diagram is that the corrosion rate of the active metal will increase as the concentration of reducible species is increased. Effect of Increasing Solution Velocity. Figure 29 shows the effect of increasing the velocity of the solution for a corroding system under cathodic control. In this instance, the corrosion rate is being controlled by the intersection of the limiting current under concentration control with the anodic polarization curve of the active metal. The effect of increasing the velocity of the solution is to decrease the thickness of the
Fig. 29
Effect of increasing the velocity of the solution for a corroding system under cathodic concentration control
93
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Corrosion: Understanding the Basics
diffusion layer adjacent to the electrode surface. As shown in the section on concentration polarization, decreasing the thickness of the diffusion layer increases the value of the limiting current and shifts it farther to the right on the diagram. The portion of the cathodic polarization curve described by activation control is not affected by increasing velocity. The net effect of increasing velocity is to increase the corrosion rate of an active metal under these conditions of concentration control for the reduction reaction. Effect of Inhibitors. Mixed-potential theory diagrams have been used to identify the mechanism of inhibition in corrosion reactions. The effect of an anodic inhibitor and a mixed inhibitor on the anodic and cathodic polarization curves and the corrosion rate of iron in H2SO4 is shown in Fig. 30 for the inhibition of corrosion of iron in H2SO4. The addition of anodic inhibitor significantly reduced the value of the anodic current at any potential; there was little effect on the cathodic polarization curve. The net effect was to shift the anodic polarization curve to the left, with a corresponding decrease in corrosion current from i1 to i2 and a corresponding increase in (more positive) corrosion potential from E1 to E2. This inhibitor is referred to as anodic inhibitor because its major influence on the corrosion reaction results from its decrease in the anodic reaction kinetics. The mixed inhibitor significantly reduced both the cathodic reaction kinetics and the anodic reaction kinetics, and both the cathodic and anodic curves were shifted to lower current values (to the left in Fig. 30). The net effect in this instance was to significantly reduce the corrosion Anodic
Mixed
Potential, V-SHE
–0.1
–0.2
E2
E2
E1
E1 –0.3
i2
i2
i1 i1
–0.4
10–5
10–4
10–3
10–2
10–5
10–4
10–3
10–2
Current density, A/cm2
Fig. 30
Effect of an anodic inhibitor and a mixed inhibitor on the anodic and cathodic polarization curves and corrosion rate of iron in sulfuric acid. Dashed lines, before addition of inhibitor; solid lines, after addition of inhibitor
Principles of Aqueous Corrosion
current from i1 to i2 and to increase the corrosion potential from E1 to E2. It should be apparent from Fig. 30 that simply measuring the effects of an inhibitor on corrosion potential does not provide a definitive picture of inhibitor behavior. The figure also shows the benefit of determination of anodic and cathodic polarization curves as a function of inhibitor type and concentration. Useful insight as to the mechanism of inhibition can be gained.
Exchange Currents The concept of exchange current is important in mixed-potential theory. The magnitude of the exchange current for a given half-cell reaction can greatly affect the resulting corrosion rate observed in an operating corrosion cell. The exchange current is defined as the steady-state value of current for a given half-cell reaction. This is demonstrated in Fig. 31 for the equilibrium or steady-state free-energy condition for the hydrogen evolution reaction. The electrochemical reaction for the reduction of hydrogen ions and evolution of hydrogen gas is shown. The reaction proceeding from left to right results in the evolution of hydrogen gas and is a reduction reaction. The reaction proceeding from right to left results in the generation of hydrogen ions and is an oxidation reaction. When the free energy of the hydrogen ions is equal to the free energy of hydrogen gas, the system is at steady state or equilibrium. At this condition, the rate at which hydrogen ions are consumed to generate hydrogen gas is equal to the rate at which hydrogen gas is consumed to generate hydrogen ions. The forward and reverse reactions occur at equal rates. The magnitude of the current where the forward and reverse reactions are equal is defined as the exchange current. The magnitude of the exchange current depends on the properties of the electrode surface upon which the reaction occurs. This is demonstrated
Fig. 31
Equilibrium or steady-state free-energy condition for hydrogen evolution
95
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Corrosion: Understanding the Basics
2H + +
2
–
+
Pb
Fe,Cu
Pt
H2 ++
2H
0
e
–
2e
EH
H2
– –12
–9
–6
–3
log i, Amp/cm2
Fig. 32
Exchange-current densities for hydrogen evolution of lead, iron, copper, and platinum
for the hydrogen evolution reaction in Fig. 32, which shows mixed-potential theory diagrams for the oxidation and reduction reactions of hydrogen evolution on lead, iron, copper, and platinum. The anodic reaction is the consumption of hydrogen gas with the production of hydrogen ions, and the reduction reaction is the consumption of hydrogen ions with the production of hydrogen gas. The steady-state value is described by the intersection of these curves. For lead, the curves intersect at an exchange-current value of 10–12 A/cm2. For iron and copper, the exchange current is 10–6 A/cm2. For platinum, the exchange current is 10–3 A/cm2. For these metals, the exchange current for hydrogen evolution varies over a range of nine orders of magnitude. Lead is the least efficient cathodic surface, and the hydrogen evolution reaction proceeds relatively slowly; platinum is the most efficient cathodic surface, and the hydrogen evolution reaction occurs much more rapidly. Each half-cell reaction has its characteristic exchange current on the particular electrode surface. Referring to Fig. 23, the exchange current for hydrogen evolution on zinc is seen to be approximately 10–10 A/cm2, and the exchange current for the zinc to zinc ion reaction is seen to be approximately 10–7 A/cm2. These values are indicated by i0 for each reaction in Fig. 23. The magnitude of the exchange current can significantly affect the corrosion rate. The effect of the exchange current for the hydrogen evolution reaction on mercury, zinc, and platinum on the corrosion rate of an active metal is shown in Fig. 33. Platinum is the most efficient cathodic surface for this reaction, followed by zinc and then mercury. On the platinum surface, the hydrogen evolution reaction proceeds much more rapidly, and the cathodic polarization curve is shifted far to the right. On mercury, the hydrogen evolution reaction occurs relatively slowly, and the curve is shifted to the left. The corrosion rate of the active metal cou-
Principles of Aqueous Corrosion
Fig. 33
97
Effect of increasing the efficiency of cathodic reaction surfaces on the corrosion rate of an active metal
pled with hydrogen evolution on a mercury, zinc, or platinum surface is indicated by i1, i2, and i3, respectively. The corrosion rate in the presence of a highly efficient cathodic surface such as platinum can be many orders of magnitude greater than for a relatively sluggish cathodic surface such as mercury.
References Selected References · · · · · · · ·
Corrosion, Vol 13, ASM Handbook, ASM International, 1987 M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill, 1986 D.A. Jones, Principles and Prevention of Corrosion, Prentice Hall, 1996 D.L. Piron, The Electrochemistry of Corrosion, NACE International, 1991 M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solutions, NACE International, 1974 J.C. Skully, The Fundamentals of Corrosion, 3rd ed., Pergamon Press, 1990 L.L Shreir, Electrochemical Principles of Corrosion, National Corrosion Service, National Physical Laboratories, Teddington, Middlesex, United Kingdon J.M Smith and H.C. Van Hess, Introduction to Chemical Engineering Thermodynamics, McGraw-Hill, 1975
Corrosion: Understanding the Basics J.R. Davis, editor, p99-192 DOI: 10.1361/cutb2000p099
CHAPTER
Copyright © 2000 ASM International® All rights reserved. www.asminternational.org
4
Forms of Corrosion: Recognition and Prevention CORROSION PROBLEMS can be divided into eight categories based on the appearance of the corrosion damage or the mechanism of attack: · Uniform or general corrosion · Pitting corrosion · Crevice corrosion, including corrosion under tubercles or deposits, filiform corrosion, and poultice corrosion · Galvanic corrosion · Erosion-corrosion, including cavitation erosion and fretting corrosion · Intergranular corrosion, including sensitization and exfoliation · Dealloying · Environmentally assisted cracking, including stress-corrosion cracking (SCC), corrosion fatigue, and hydrogen damage (including hydrogen embrittlement, hydrogen-induced blistering, high-temperature hydrogen attack, and hydride formation)
Although these forms are presented in the context of aqueous corrosion, many of them are also operative at high temperature. For example, high-temperature corrosion by oxidation or sulfidation can take the form of uniform attack, pitting, or dealloying. Molten metal or molten
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Corrosion: Understanding the Basics
salt environments can produce uniform corrosion, dealloying, or intergranular attack. While the corrosion classification scheme listed above is convenient, it should be emphasized that it is arbitrary and by no means perfect. Many corrosion problems are due to more than one form of corrosion acting simultaneously. For example, pitting corrosion may be caused by crevice corrosion, deposit corrosion, cavitation, or fretting corrosion. Additionally, in some metal systems where dealloying may occur, this form of corrosion may be a precursor to stress-corrosion cracking. Similarly, deep pits can act as stress raisers and serve as nucleation sites for corrosion fatigue failures. Some forms of corrosion such as stray-current corrosion and deposition corrosion are extremely difficult to classify. Stray-current corrosion is different from natural corrosion because it is caused by an extremely induced electrical current and is basically independent of some of the environmental factors that influence other forms of corrosion. Deposition corrosion is a combination of pitting and galvanic corrosion that can occur in a liquid environment when ions of more cathodic metal (e.g., copper) are plated out of solution onto a more anodic metal surface (e.g., aluminum). Despite its shortcomings, the classification of corrosion forms based on physical appearance or attack mechanism allows a large and complex technology to be broken down into more usable and understandable pieces. This classification system is particularly useful for failure analysis, guiding investigators in the determination of contributing factors and of methods for controlling the specific form of corrosion.
Uniform Corrosion General Description. Uniform or general corrosion, as the name implies, results in a fairly uniform penetration (or thinning) over the entire exposed metal surface. The general attack results from local corrosioncell action; that is, multiple anodes and cathodes are operating on the metal surface at any given time. The location of the anodic and cathodic areas continues to move about on the surface, resulting in uniform corrosion. Uniform corrosion represents the greatest destruction of metal on a tonnage basis. This form of corrosion however, is not of too great concern from a technical standpoint because the life of equipment can be accurately estimated on the basis of comparatively simple immersion tests. These tests allow weight (mass) loss to be monitored, and the reduction of thickness as a function of time can be calculated. Corrosion rate expressions and allowances for general corrosion are described in Chapter 2.
Forms of Corrosion: Recognition and Prevention
Uniform corrosion often results from atmospheric exposure (especially polluted industrial environments); exposure in fresh, brackish, and salt waters; or exposure in soils and chemicals. Corrosion in these environments is discussed in Chapter 5. Metals Affected. All metals are affected by uniform corrosion, although passive materials, such as stainless steels or nickel-chromium alloys are normally subjected to localized forms of attack. The rusting of steel (Fig. 1), the green patina formation on copper, and the tarnishing of silver are typical examples of uniform corrosion. In some metals, such as steel, uniform corrosion produces a somewhat rough surface by removing a substantial amount of metal, which either dissolves in the environment or reacts with it to produce a loosely adherent, porous coating of corrosion products. In such reactions as in the tarnishing of silver in air, the oxidation of aluminum in air, or attack on lead in sulfatecontaining environments, thin, tightly adherent protective films are produced, and the metal surface remains smooth. Prevention. Uniform corrosion can be prevented or reduced by proper materials selection, the use of coatings or inhibitors, or cathodic or anodic protection. These corrosion prevention methods can be used individually or in combination. Uniform corrosion is often treated by building a corrosion allowance into the structure. If the corrosion rate is
Fig. 1
Uniform corrosion (rusting) of a weathering steel highway bridge girder
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Corrosion: Understanding the Basics
100 mm/year, then the addition of 500 mm to the thickness of metal will provide 5 years of operation. The measure of uniform corrosion rate required for determination of the corrosion allowance is estimated from prior conditions in similar service, data presented in the corrosion literature, and experimental data determined from coupon exposures. More detailed information on prevention of uniform corrosion can be found in Chapters 7 to 10. Corrosion allowances are also discussed in Chapter 2.
Pitting Corrosion General Description. Pitting is a highly localized form of corrosion that produces sharply defined holes. These holes may be small or large in diameter, but in most cases, they are relatively small. Pits may be isolated from each other on the surface or so close together that they resemble a roughened surface. Variations in the cross-sectional shape of pits are shown in Fig. 2. Every engineering metal or alloy is susceptible to pitting. Pitting occurs when one area of a metal becomes anodic with respect to the rest of the surface or when highly localized changes in the corrodent in contact with the metal, as in crevices, cause accelerated localized attack. Difficulty of Detection. Pitting is one of the most insidious forms of corrosion. It can cause failure by perforation while producing only a small weight loss on the metal. Also, pits are generally small and often remain undetected. A small number of isolated pits on a generally uncorroded surface are easily overlooked. A large number of very small pits on a generally uncorroded surface may not be detected by simple visual examination, or their potential for damage may be underestimated.
Fig. 2
Variations in the cross-sectional shape of pits. (a) Narrow and deep. (b) Elliptical. (c) Wide and shallow. (d) Subsurface. (e) Undercutting. (f) Shapes determined by microstructural orientation. Source: ASTM G 46
Forms of Corrosion: Recognition and Prevention
When pits are accompanied by slight or moderate general corrosion, the corrosion products often mask them. Pitting is sometimes difficult to detect in laboratory tests and in service because there may be a period of months or years, depending on the metal and the corrodent, before the pits initiate and develop to a readily visible size. Delayed pitting sometimes occurs after an unpredictable period of time in service, when some change in the environment causes local destruction of a passive film. When this occurs on stainless steels, for example, there is a substantial increase in solution potential of the active area, and pitting progresses rapidly. Stages of Pitting. Immediately after a pit has initiated, the local environment and any surface films on the pit-initiation site are unstable, and the pit may become inactive after just a few minutes if convection currents sweep away the locally high concentration of hydrogen ions, chloride ions, or other ions that initiated the local attack. Accordingly, the continued development of pits is favored in a stagnant solution. When a pit has reached a stable stage, barring drastic changes in the environment, it penetrates the metal at an ever-increasing rate by an autocatalytic process. In the pitting of a metal by an aerated sodium chloride solution, rapid dissolution occurs within the pit, while reduction of oxygen takes place on adjacent surfaces. This process is selfpropagating. The rapid dissolution of metal within the pit produces an excess of positive charges in this area, causing migration of chloride ions into the pit (Fig. 3). Thus, in the pit there is a high concentration of metal chlorides (M+Cl–) and as a result of hydrolysis, a high concentration of hydrogen ions. Both hydrogen and chloride ions stimulate the dissolution of most metals and alloys, and the entire process accelerates with time. Because the solubility of oxygen is virtually zero in concentrated solutions, no reduction of oxygen occurs within a pit. Cathodic reduction of oxygen on the surface areas adjacent to pits tends to suppress corrosion on these surface areas. Thus, isolated pits cathodically protect the surrounding metal surface. Because the dense, concentrated solution within a pit is necessary for its continuing development, pits are most stable when growing in the direction of gravity. Also, the active anions are more easily retained on the upper surfaces of a piece of metal immersed in or covered by a liquid. Some causes of pitting are local inhomogeneity on the metal surface, local loss of passivity, mechanical or chemical rupture of a protective oxide coating, galvanic corrosion from a relatively distant cathode, the formation of a metal ion or oxygen concentration cell under a solid deposit (crevice corrosion), and the presence of biological organisms. The rate of pitting is related to the aggressiveness of the corrodent at the site of pitting and the electrical conductivity of the solution containing the corrodent. For a given metal, certain specific ions increase the
103
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Corrosion: Understanding the Basics
probability of attack from pitting and accelerate that attack once initiated. Pitting is usually associated with metal-environment combinations in which the general corrosion rate is relatively low. For a given combination, the rate of penetration into the metal by pitting can be 10 to 100 times that by general corrosion. With carbon and low-alloy steels in relatively mild corrodents, pits are often generally distributed over the surface and change locations as they propagate. If they blend together, the individual pits become virtually indistinguishable, and the final effect is a roughened surface but a generally uniform reduction in cross section. If the initial pits on carbon steel do not combine in this way, the result is rapid penetration of the metal at the sites of the pits and little general corrosion.
Fig. 3
Autocatalytic processes occurring in a corrosion pit. The metal, M, is being pitted by an aerated sodium chloride (NaCl) solution. Rapid dissolution occurs within the pit, while oxygen reduction takes place on the adjacent surfaces. A more detailed explanation of this self-sustaining process is given in Ref 1.
Forms of Corrosion: Recognition and Prevention
The most common causes of pitting in steels are surface deposits that set up local concentration cells and dissolved halides that produce local anodes by rupture of the protective oxide film. Anodic corrosion inhibitors, such as chromates, can cause rapid pitting if present in concentrations below a minimum value that depends on the metal-environment combination, temperature, and other factors. Pitting also occurs at mechanical ruptures in protective organic coatings if the external environment is aggressive or if a galvanic cell is active. With corrosion-resistant alloys, such as stainless steels, the most common cause of pitting corrosion is highly localized destruction of passivity by contact with moisture that contains halide ions, particularly chlorides. Figure 4 shows deep pits that formed in a type 316 stainless steel centrifuge head from a calcium chloride (CaCl2) solution. Chloride-induced pitting of stainless steels usually results in undercutting (see Fig. 2e), producing enlarged subsurface cavities or caverns. Undercutting also occurs when most metals are exposed to highly acidic conditions (Fig. 5). Pitting of Various Metals. Pitting occurs in most commonly used metals and alloys. Iron buried in the soil corrodes with the formation of shallow pits, but carbon steels in contact with hydrochloric acid or stainless steels immersed in seawater characteristically corrode with the formation of deep pits (Fig. 4 and 6). Aluminum tends to pit in waters containing chloride ions (for example, at stagnant areas), and aluminum brasses are subject to pitting in polluted waters. Despite their good resistance to general corrosion, stainless steels are more susceptible to pitting than many other metals. High-alloy stainless steels containing chromium, nickel, and molybdenum are also more resistant to pitting but are not immune under all service conditions. Pitting failures of corrosion-resistant alloys, such as Hastelloy C, Hastelloy G, and Incoloy 825, are relatively uncommon in solutions that do not contain halides, although any mechanism that permits the establishment of an electrolytic cell in which a small anode is in contact with a large cathodic area offers the opportunity for pitting attack. Prevention. Typical approaches to alleviating or minimizing pitting corrosion include the following: · Reduce the aggressiveness of the environment, for example, chloride ions concentration, temperature, acidity, and oxidizing agents · Upgrade the materials of construction, for example, use molybdenumcontaining (4 to 6% Mo) stainless steels, molybdenum + tungsten nickel-base alloys, overalloy welds, and use corrosion-resistant alloy linings · Modify the design of the system, for example, avoid crevices and the formation of deposits, circulate/stir to eliminate stagnant solutions, and ensure proper drainage
105
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Corrosion: Understanding the Basics
Fig. 4
Two views of deep pits in a type 316 stainless steel centrifuge head due to exposure to CaCl2 solution
Forms of Corrosion: Recognition and Prevention
107
(a)
(b)
Fig. 5
(a) Pitting of a carbon steel pipe exposed to a strong mineral acid. (b) Close-up view shows the narrow pit mouths and the pronounced undercutting. Source: Nalco Chemical Company
Crevice Corrosion General Description. Crevice corrosion is a form of localized attack that occurs at narrow openings or spaces (gaps) between metal-to-metal or nonmetal-to-metal components. This type of attack results from a concentration cell formed between the electrolyte within the crevice, which is oxygen starved, and the electrolyte outside the crevice, where oxygen is more plentiful. The material within the crevice acts as the anode, and the exterior material becomes the cathode. Crevices may be produced by design or accident. Crevices caused by design occur at gaskets, flanges, rubber O-rings, washers, bolt holes, rolled tube ends, threaded joints, riveted seams, overlapping screen
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Corrosion: Understanding the Basics
wires, lap joints, beneath coatings (filiform corrosion) or insulation (poultice corrosion), and anywhere close-fitting surfaces are present. Occluded regions are also formed under tubercles (tuberculation), deposits (deposit corrosion), and below accumulations or biological materials (biologically influenced corrosion). Similarly, unintentional crevices such as cracks, seams, and other metallurgical defects could serve as sites for corrosion. Resistance to crevice corrosion can vary from one alloy-environment system to another. Although crevice corrosion affects both active and passive metals, the attack is often more severe for passive alloys, particularly those in the stainless steel group. Breakdown of the passive film within a restricted geometry leads to rapid metal loss and penetration of the metal in that area. Numerous interrelated metallurgical, geometrical, and environmental factors, as well as electrochemical reactions, affect both crevice initiation and propagation. A number of these factors are indicated in Table 1. Crevice Corrosion Propagation. The propagation of crevice corrosion is thought to involve the dissolution of metal and the maintenance of a high degree of acidity within the crevice solution by hydrolysis of the dissolved metal ions (Ref 1). The crevice corrosion propagation process is illustrated schematically in Fig. 7 for stainless steel corroding in a
Fig. 6
Deep pits on a carbon steel check valve that was inadvertently exposed to hydrochloric acid during a plant upset. Note how pits intersect to form areas of jagged metal loss. A steel probe tip is also shown in the photo. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
109
neutral aerated sodium chloride solution. The anodic metal dissolution reaction within the crevice, M ® Mn+ + ne–, is balanced by the cathodic reaction on the adjacent surface, O2 + 2H2O + 4e– ® 4OH–. The increased concentration of M+ within the crevice results in the influx of chloride ions (Cl–) to maintain neutrality. The metal chloride formed, M+Cl–, is hydrolyzed by water to the hydroxide and free acid as: M+Cl– + H2O ® MOH + H+Cl– The acid produced by the hydrolysis reaction keeps the pH to values below 2 (Ref 2), while the pH of the solution outside the crevice remains neutral (pH 7). In simple terms, the electrolyte present within an actively corroding crevice can be regarded as concentrated hydrochloric acid containing metal chlorides dissolved at concentrations near saturation. Examples of Crevice Corrosion. Figure 8 shows crevice corrosion of a type 304 stainless steel fastener removed from a seawater jetty after 8 years. Although the washer shows severe deterioration, the function Table 1 Factors that can affect the crevice corrosion resistance of various alloys Factor
Type
Geometrical
Type of crevice Metal-to-metal Nonmetal to metal Crevice gap (tightness) Crevice depth Exterior-to-interior surface area ratio Bulk solution O2 content pH Chloride level Temperature Agitation Mass transport, migration Diffusion and convection Crevice solution: hydrolysis equilibria Biological influences Metal dissolution O2 reduction H2 evolution Alloy composition Major elements Minor elements Impurities
Environmental
Electrochemical reactions
Metallurgical
O2
OH–
Cl–
Crevice M2+ H+
e–
Fig. 7
A schematic of the crevice corrosion propagation mechanism
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Corrosion: Understanding the Basics
Fig. 8
Crevice corrosion at a metal-to-metal crevice site formed between components of type 304 stainless steel fastener in seawater
Fig. 9
Crevice corrosion at nonmetal gasket site on an alloy 825 seawater heat exchanger
Forms of Corrosion: Recognition and Prevention
111
of the fastener was not diminished. On the other hand, Fig. 9 shows crevice corrosion beneath the water-box gasket of an alloy 825 (44Ni-22Cr-3Mo-2Cu) seawater heat exchanger that allowed sufficient leakage to warrant shutdown and replacement after only 6 months. In cases in which the bulk environment is particularly aggressive, general corrosion may preclude localized corrosion at a crevice site. Figure 10 compares the behavior of type 304 and type 316 stainless steels exposed in different zones of a model sulfur dioxide (SO2) scrubber. In the aggressive acid condensate zone, type 304 incurred severe general corrosion of the exposed surfaces, while the more resistant type 316 suffered attack beneath a polytetrafluoroethylene (PTFE) insulating spacer. In the higher pH environment of the limestone slurry zone, type 304 was resistant to general corrosion but was susceptible to crev-
(a)
(b)
(c)
Fig. 10
Variation in stainless steel corrosion resistance in model SO2 scrubber environments. (a) Type 304 in acid condensate. (b) Type 316 in acid condensate. (c) Type 304 in limestone slurry zone
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Corrosion: Understanding the Basics
ice corrosion. Other alloy systems, such as aluminum and titanium, may also be susceptible to crevice corrosion. For aluminum, the occurrence of crevice corrosion would depend on the passivity of the particular alloy. In most cases, general corrosion would likely preclude crevice corrosion. Titanium alloys are typically quite resistant but may be susceptible to crevice corrosion in elevated temperature, chloride-containing acidic environments. Although the occurrence of crevice corrosion in cast irons and carbon steels is not frequent, the presence of chlorides and/or crevices or other shielding areas presents conditions that can be favorable to crevice attack. Rust often accumulates at crevice mouths, and darker oxides are present within the crevices. Figure 11 shows a cast-iron valve block that exhibited crevice corrosion beneath rubber O-rings. In seawater, localized corrosion of copper and its alloys at crevices is different from that of stainless-type materials because the attack occurs outside of the crevice rather than within. In general, the degree of crevice- related attack increases. Figure 12 compares the crevice corrosion behavior for several different materials exposed to ambient-temperature seawater for various periods. In each case, a nonmetallic washer created the crevice. The more classical form of crevice corrosion (that is, beneath the crevice former) is shown for type 904L stainless steel (20Cr-25Ni-4.5Mo-1.5Cu) after only 30 days of exposure (Fig. 12a). For 70Cu-30Ni, corrosion occurred just outside of the crevice mouth and was found to be quite shallow after 6 months (Fig. 12b). In contrast, crevice-related corrosion of alloy 400 (70Ni-30Cu) was more severe after only 45 days (Fig. 12c). In some cases, corrosion may occur within as well as outside of the crevice.
Fig. 11
Closeup of annular regions below rubber O-rings on a cast-iron valve block. Note how damage varies from hole to hole, probably due to variation in the crevice geometry. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
(a)
(b)
(c)
Fig. 12
Crevice-related corrosion for different alloys in natural seawater. (a) Alloy 904L (20Cr-25Ni-4.5Mo-1.5Cu) after 30 days. (b) 70Cu-30Ni after 180 days. (c) Alloy 400 (70Ni-30Cu) after 45 days
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Prevention. Many of the factors listed in Table 1 must be considered if crevice corrosion is to be eliminated or minimized. Wherever possible, crevices should be eliminated at the design stage (guidelines and illustrations are presented in Chapter 7). When unavoidable, they should be kept as open and shallow as possible to allow continued entry of the bulk environment. As with pitting corrosion, a common way to protect against crevice corrosion is to select more corrosion-resistant materials (e.g., molybdenumcontaining, more highly alloyed stainless steels and nickel-base alloys). Weld overlays can also be employed. The influence of gasket material on nonmetal-to-stainless steel crevice formation should also be considered. Natural and synthetic elastomertype gaskets are less likely to promote crevice corrosion than polytetrafluoroethylene (PTFE) gaskets (with or without glass fiber) and para-aramid fiber + nitrile binder-type gaskets. While carbon and graphite-containing gaskets promote crevice, PTFE and para-aramid + nitrile promote even greater attack. Cleanliness is an important factor, particularly when conditions promote deposition on metal surfaces. Regular cleaning involving mechanical methods is commonly employed. Filters can also be employed to remove materials that can deposit on the metal surface. It is very important to try to avoid using hydrochloric acid to clean stainless steel systems. Chloride will concentrate in preexisting crevices during cleaning and may not be removed subsequently.
Tuberculation (Ref 3) General Description. Tuberculation, which is a specific type of crevice corrosion often encountered in cooling water systems, is defined as “the formation of localized corrosion products scattered over the surface in the form of knoblike mounds called tubercles.” Tubercles can choke pipes, leading to diminished flow and increased pumping costs (Fig. 13). Tubercles form on steel and cast iron when surfaces are exposed to oxygenated waters. Soft waters with high bicarbonate alkalinity stimulate tubercle formation, as do high concentrations of sulfate, chloride, and other aggressive anions (Ref 3). The formation of tubercles by biological organisms acting in conjunction with electrochemical corrosion also occurs in many aqueous environments. Sulfatereducing and acid-producing bacteria associated with biological organisms accelerate crevice attack. It should be noted, however, that tubercles frequently form without the presence of any biological organisms, and the following discussion does not take into account biological effects. Chapter 5 describes the influence of biological organisms and biofilms on corrosion.
Forms of Corrosion: Recognition and Prevention
Features and Growth Characteristics. Tubercles are much more than amorphous lumps of corrosion product and deposit. As shown in Fig. 14(a), they are highly structured and consist of five distinct layers (Ref 4): · Outer crust (primarily red, brown, and orange corrosion products (i.e., rust) · Inner shell (magnetite) · Core material (ferrous hydroxide) · Fluid-filled cavity (containing Fe2+, Cl–, and SO 24 + ) · Corroded floor, which is almost always a dish-shaped depression that is much wider than it is deep (Fig. 15). Average corrosion rates are usually 0.5 mm/year (20 mils/year) or less.
Typical reactions occurring within these five layers are shown in Fig. 14(b). As rust accumulates, oxygen migration is reduced through the corrosion product layer. Regions below the rust layer become oxygen depleted. An oxygen concentration cell then develops. Corrosion naturally becomes concentrated into small regions beneath the rust, and volcanolike structures, or tubercles, are generated.
Fig. 13
Heavily tuberculated 75 mm (3 in.) outer diameter steel mill water supply line. Source: Nalco Chemical Company
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The unique features of tubercles can better be appreciated by examining the series of photos in Fig. 16 and 17 that show a tuberculated 1010 carbon steel baffle plate from a test water box in a utility cooling system. After a two-week period to monitor corrosion and fouling, small, hollow incipient tubercles dotted the surface of the plate (Fig. 16a). Each tubercle capped a depression (pit) no deeper than 0.13 mm (0.005 in.) (Fig. 16b). This indicated local average corrosion rates were as high as 3.3 mm/year (130 mils/year). Each tubercle exhibited small clam-shell marks or growth rings (Fig. 17a). Each ring was formed by fracture at the tubercle base during growth. Ejected internal contents rapidly deposited Crust (friable) • Hematite–red, brown, orange (ferric hydroxide) • Carbonate–white • Silicates–white Fluid-filled cavity (Fe++, Cl–, SO4=)
Shell (brittle) • Magnetite–black
Water
Core (friable) • Ferrous hydroxide– greenish-black • Iron carbonate– gray-black (siderite) • Phosphates, etc.
Fracture in crust Corroding floor
Metal loss region Metal (a) Fe(OH)2 + 1/2H2O + 1/4 O2 Cathode 2e– + H2O + 1/2 O2
2OH–
OH–
e–
OH–
OH–
Fe++ + CO3= FeCO3 Fe++ + 2OH– Fe(OH)2
OH–
Fe++
OH– CO3= HCO3 Cl– SO4–
Fe(OH)3
Fe++
Fe++Cl2– + 2H2O
Fe
OH–
Fe(OH)2 + 2H+ Cl– Fe++
2H+Cl– + Fe++
Migration of negative ion into tubercle
Fe++Cl2– + 2H+
Fe++ e–
Anode Fe++ + 2e–
(b)
Fig. 14
Schematics of tubercles formed on iron or steel in oxygenated waters. (a) Structural features and associated compounds. (b) Corrosion reactions within the tubercle. Source: Ref 4
Forms of Corrosion: Recognition and Prevention
when contacting oxygenating water. Tubercles were hollow (Fig. 17b), and the surfaces below the cap contained concentrations exceeding 10% of chloride and sulfate, producing severe localized acidic conditions. Prevention. Tuberculation can be prevented or minimized by the following: · Using inhibitors · Altering system operation (e.g., controlling water flow and temperature conditions) · Coating the carbon steel or cast iron components with epoxy or other field-applied or factory-applied organic coatings · Using more corrosion-resistant materials such as stainless (not sensitized), brasses, copper-nickels, titanium, or aluminum. None of these materials will form tubercles in oxygenated water. However, each of these alloys may suffer other problems that would preclude their use in a specific environment.
Fig. 15
Perforation at a dish-shaped depression on the internal surface of a large-diameter steel pipe. A large tubercle capped the depression but was dislodged during tube sectioning. Source: Nalco Chemical Company
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Deposit Corrosion General Description. Deposit corrosion refers to crevice corrosion occurring under naturally occurring nonman-made deposits. It has also been referred to as “underdeposit corrosion” in the literature (Ref 3).
(a)
(b)
Fig. 16
Tuberculation of a steel baffle plate. (a) Numerous incipient tubercles formed in 2 weeks. (b) Tubercles removed to show pitlike depressions beneath each mound. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
119
Deposits include water-borne precipitates (e.g., carbonates, silicates, and phosphates in cooling water systems), transported particulate, corrosion products (e.g., manganese-rich deposits or iron oxides), biological materials, and a variety of contaminants such as grease, oil, process chemicals, silt, sand, and road debris (e.g., salt, mud, and water deposited
(a)
40 mm
(b)
100 mm
Fig. 17
Scanning electron micrographs of the tubercles shown in Fig. 16. (a) Clam-shell growth steps formed by successive fractures of the tubercle base. (b) Tubercles broken open to reveal hollow interiors. Source: Nalco Chemical Company
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on the underside of automobile fenders). Associated corrosion is fundamentally related to whether deposits are innately aggressive (e.g., deposits containing corrosive substances such as sulfur-containing and chlorine-containing species) or simply serve as an occluding medium beneath which concentration cells develop. Severe wastage from pitting or general corrosion can occur under these deposits. Deposits that are heavily stratified can also clog pipelines (Fig. 18). Figure 19 shows deposits containing organic acids formed by oxidation of rolling mill oils. Up to 40% by weight of the lumps shown in Fig. 19 are iron oxides, hydroxides, and organic-acid iron salts. Metals Affected. Unlike tuberculation, which is associated with only cast irons and steels, deposit corrosion affects a wide variety of metals and is of particular concern with passive alloys such as stainless steels, aluminum, nickel, and titanium. Figure 20 shows how the concentration of aggressive ions beneath deposits can produce severe localized corrosion on stainless steels. Materials selection should be based on careful testing in the specific service environment anticipated. Prevention. Deposit-related corrosion may be minimized by the following: · Regular cleaning to remove deposits
Fig. 18
Thick calcium carbonate deposits on a condenser tube and a copper transfer pipe. Heavily stratified deposits reflect changes in water chemistry, heat transfer, and flow. Corrosion may be slight beneath heavy accumulations of fairly pure calcium carbonate because such layers can inhibit some forms of corrosion. However, calcium carbonates are often intermixed with silt, metal oxides, and other precipitates, leading to severe deposit attack. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
· Design changes. Deposition caused by settling of suspended particulate may be reduced by increasing flow. Dead legs, stagnant areas, and other low-flow regions should be eliminated if possible. · Water treatments, such as removing suspended solids, the use of biodispersants and biocides in biofouled systems, and the judicious use of inhibitors · Cathodic protection
Fig. 19
Carbon steel coupon removed from a rolling mill cooling tank. Note the thick greasy deposits resulting from rolling oils. Removal of the deposits shows the corrosion beneath (see right side of figure). Source: Nalco Chemical Company
Fig. 20
Severe localized corrosion on a type 316 stainless steel heat exchanger tube. Attack occurred beneath deposits, which were removed to show wastage. Source: Nalco Chemical Company
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· Protective coatings such as water-impermeable phenolic resins, epoxies, and other organic compounds, sacrificial coatings (e.g., zinc thermal spray coatings), and electroplated coatings
Filiform Corrosion General Description. Filiform corrosion is a special case of concentration-cell corrosion that occurs on metallic surfaces coated with a thin organic film that is typically 0.1 mm (4 mils) thick. The pattern of attack is characterized by the appearance of fine filaments emanating from one or more sources in semirandom directions. The source of initiation is usually a defect or scratch in the coating. The filaments are fine tunnels composed of corrosion products underneath the bulged and cracked coating. Filiforms are visible at an arm’s length as small blemishes. Upon closer examination, they appear as fine striations shaped like tentacles or cobweblike traces (Fig. 21). A filiform has an active head and a filamentous tail (Fig. 22). Filiform attack occurs when the relative humidity is typically between 65 and 90% in most cases. The average width of a filament varies between 0.05 to 3 mm (2 to 120 mils). Filament width depends on the
Fig. 21
A lacquered steel can lid exhibiting filiform corrosion showing both large and small filaments partially oriented in the rolling direction of the steel sheet. Without this 10´ magnification by a light microscope, the filiforms look like fine striations or minute tentacles.
Forms of Corrosion: Recognition and Prevention
123
coating, the ambient relative humidity, and the corrosive environment. Typical filament height is about 20 mm (0.8 mil). The filament growth rate can also vary widely, with rates observed as low as 0.01 mm/day (0.4 mil/day) and up to a maximum rate of 0.85 mm/day (35 mils/day). The depth of attack in the filiform tunnels can be as deep as 15 mm (0.6 mil). The fluid in the leading head of a filiform is typically acidic, with a
(a)
(b)
(c)
Fig. 22
Filiform corrosion of PVC-coated aluminum foil. (a) Advancing head and cracked tail section of a filiform cell. Scanning electron microscopy (SEM), 80´. (b) The gelatinous corrosion products of aluminum oozing out of the porous end tail section of a filiform cell. SEM. 830´. (c) Tail region of a filiform cell. Tail appears iridescent due to internal reflection. Light microscopy, 60´
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pH from 1 to 4. In all cases, oxygen or air and water were needed to sustain filiform corrosion. Metals Affected. Filiform corrosion routinely occurs on coated steel cans, coated tin-plated steel, aluminum foil-laminated packaging, painted aluminum, painted magnesium, and other lacquered metallic items placed in areas subjected to high humidity. Growth rates for filiform corrosion on various coated metals are listed in Table 2. Prevention. Prevention of filiform corrosion can be accomplished by the following: · Reduction of relative humidity to below 60% by the use of drying fans, dehumidifiers, or the addition of desiccants in packaging applications · Use of zinc (galvanizing) and zinc primers on steels · Use of zinc chromate primers, chromic acid anodizing, and chromate or chromate-phosphate conversion coatings on aluminum · Use of multiple-coat paint systems · Use of less active metal substrates (e.g., copper, stainless steel, or titanium)
Table 2
Filiform corrosion growth rates on various coated metals
Coating
Initiating environment
Typical rate mm/day mils/day
Relative humidity, %
Filament width mm mils
Steels Varnish Copol Lacquer Linseed oil Alkyds Alkyd urea Epoxy urea Epoxy Acrylic Polyurethane Polyester
NaCl Acetic acid NaCl Acetic acid NaCl NaCl Acetic acid FeCl2 NaCl/FeCl2 Acetic acid NaCl Acetic acid NaCl Acetic acid Acetic acid
0.33–0.53 0.5 0.03 0.85 0.04–0.08 0.50 0.1 0.26–0.43 0.01–0.46 0.16 0.19–0.86 0.1 0.16–0.50 0.9 0.08
13–21 20 1.2 33.5 1.6–3.1 20 4 10–17 0.4–18 6.3 7.5–34 4 6.3–20 3.5 3.1
65–85 86 60–94 ¼ ¼ 80 85 80 80 85 80 85 90 85 85
0.1–0.3 0.15 ¼ 0.1–0.5 0.05–0.1 0.1–0.5 ¼ 0.25 0.25 ¼ 0.25 ¼ 0.1–0.3 ¼ ¼
4–12 6 ¼ 4–20 2–4 4–20 ¼ 10 10 ¼ 10 ¼ 4–12 ¼ ¼
HCl vapor HCl vapor HCl vapor HCl vapor HCl vapor
0.1 0.1 0.1 0.2 0.09
4 4 4 4 3.5
85 85 75–85 85 85
0.5–1.0 0.5–1.0 0.5–1.0 0.5–1.0 0.5–1.0
20–40 20–40 20–40 20–40 20–40
HCl vapor HCl vapor HCl vapor HCl vapor HCl vapor
0.2 0.3 0.3 0.2 0.3
8 12 12 8 12
75 75 75 75 75
¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ ¼ ¼
Aluminum alloys Alkyds Acrylic Polyurethane Polyester Epoxy Magnesium Alkyds Acrylic Polyurethane Polyester Epoxy
Forms of Corrosion: Recognition and Prevention
125
Poultice Corrosion (Ref 5) General Description. Poultice corrosion is a special case of localized corrosion due to differential aeration, which usually takes the form of pitting when an absorptive material such as paper, wood, asbestos, sacking, or cloth is in contact with a metal surface that becomes wetted periodically. No action occurs while the entire assembly is wet, but during the drying period, adjacent wet and dry areas develop. Near the edges of the wet zones, differential aeration develops, which leads to pitting, as in the case of crevice corrosion. Prevention. Poultice corrosion is prevented by avoiding the contact of absorptive materials with a metal surface, by painting the surface that will contact such materials, or by designing to prevent such materials from becoming wet in service.
Galvanic Corrosion General Description Galvanic corrosion occurs when a metal or alloy is electrically coupled to another metal or conducting nonmetal in the same electrolyte. The three essential components are the following: · Materials possessing different surface potential · A common electrolyte · A common electrical path
A mixed metal system in a common electrolyte that is electrically isolated will not experience galvanic corrosion, regardless of the proximity of the metals or their relative potential or size. During galvanic coupling, corrosion of the less corrosion-resistant metal increases, and the surface becomes anodic, while corrosion of the more corrosion-resistant metal decreases, and the surface becomes cathodic. The driving force for corrosion or current flow is the potential developed between the dissimilar metals. The extent of accelerated corrosion resulting from galvanic coupling is affected by the following factors: · · · ·
The potential difference between the metals or alloys The nature of the environment The polarization behavior of the metals or alloys The geometric relationship of the component metals or alloys
The differences in potential between dissimilar metals or alloys cause electron flow between them when they are electrically coupled in a
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Corrosion: Understanding the Basics
conductive solution. The direction of flow and, therefore, the galvanic behavior depend on which metal or alloy is more active. Thus, the more active metal or alloy becomes anodic, and the more noble metal or alloy becomes cathodic in the couple. The driving force for galvanic corrosion is the difference in potential between the component metals or alloys.
Galvanic Series When all that is necessary to know is which of the materials in a system are possible candidates for galvanically accelerated corrosion and which will be unaffected or protected, information obtained from a galvanic series in the appropriate medium is useful. A galvanic series is a list of freely corroding potentials of the materials of interest in the environment of interest, arranged in order of potential. A galvanic series is easy to use and is often all that is required to answer a simple galvanic corrosion question. The material with the most negative, or anodic, corrosion potential has a tendency to suffer accelerated corrosion when electrically connected to a material with a more positive, or cathodic (noble), potential. The disadvantages of using a galvanic series include the following: · No information is available on the rate of corrosion. · Active-passive metals may display two widely differing potentials. · Small changes in the electrolyte can cause significant changes in the potentials. · Potentials may be time dependent.
Creating a galvanic series is a matter of measuring the corrosion potentials of various materials of interest in the electrolyte of interest against a reference electrode half cell, such as saturated calomel. Preparation of a valid galvanic series for specific materials in a particular service environment must account for all the factors that affect the potential of these materials in that environment. These factors include material composition, heat treatment, surface preparation (mill scale, coatings surface finish, and so on), surface oxides occurring in air, environmental composition (trace contaminants, dissolved gases, and so on), temperature, and flow rate. Exposure time is also important, particularly for materials that form corrosion product layers. All of the precautions and warnings regarding the generation and use of a galvanic series are given in ASTM standard G 82 “Standard Guide for Development and Use of a Galvanic Series for Predicting Galvanic Corrosion Performance.” The galvanic series for metals in seawater at room temperature is presented in Fig. 23. While this galvanic series is quite useful, its use for corrosion applications in other environments or at other temperatures is
Forms of Corrosion: Recognition and Prevention
Fig. 23
Galvanic series for seawater. Dark boxes indicate active behavior of activepassive alloys.
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Corrosion: Understanding the Basics
Potential vs calomel electrode, V
inappropriate and dangerous. A galvanic series in the appropriate environment is required. Generally, the more corrosion-resistant metals, such as platinum, titanium, and silver, have the most positive potentials in seawater and are located near the upper end of the series. The more electronegative metals, such as magnesium, zinc, and aluminum, are located at the lower end of the series. When electrically coupled in a common solution, the more negative (more active) metal will be the anode of the galvanic corrosion cell, and its corrosion rate will increase. The more positive (more noble) metal will be the cathode, and its corrosion rate will decrease. Active-passive metals, such as stainless steels, can exhibit either of two widely different potentials. In the active condition, these metals have a more negative potential, and in the passive state, they have a more positive potential. In a galvanic series, these active-passive metals will be listed at two levels: one more negative for the active state and the other more positive for the passive state. Example 1: How the Galvanic Series Can Mislead (Ref 6). Figure 24 shows the potential of nickel-200, type 304 stainless steel, and a 70-30 copper-nickel alloy measured separately (i.e., not in a coupled system) against a saturated calomel reference electrode. The metals were immersed in seawater for more than 15 months. The nickel and cupronickel were relatively constant during the test. Type 304 stainless steel, on the other hand, changed its potential from above the other metals to between them, and then below, and between again. The reason for this behavior was that the stainless steel changed from a passive condition to an active one, whereby localized corrosion occurred, and then changed back to a passive condition. Therefore, alloys such as type 304 can occupy two positions in a galvanic series (refer to Fig. 23). This behavior illustrates why a simple galvanic series provides limited—and sometimes incorrect—material selection and design information. Slight variations in service conditions or in the metals themselves can cause significant changes in the relative positions of the metals in the galvanic series. +0.1 Nickel
0 Type 304 0.1 Cupronickel 0.2 0.3 0.4
0
5
10
15
Time, months
Fig. 24
Potential of metals immersed in seawater for more than 15 months. Source: Ref 6
Forms of Corrosion: Recognition and Prevention
Polarization The potential generated by a galvanic cell consisting of dissimilar metals often changes with time. This potential causes a flow of current and corrosion to occur at the anodic area—the amount of corrosion being directly proportional to the current flow. As corrosion progresses, reaction products or corrosion products may accumulate at either the anode, the cathode, or both. This accumulation reduces the rate at which corrosion proceeds. The potential of the anode drifts toward that of the cathode and vise versa. The change in potentials is called polarization—anodic polarization at the anode and cathodic polarization at the cathode. Polarization is defined as the displacement of electrode potential resulting from the effects of current flow. In most corrosion reactions, cathodic polarization is more predominant. Because the degree of cathodic polarization and its effectiveness varies with metals and alloys, something about their polarization characteristics must be known before the extent or degree of galvanic corrosion for a given couple can be predicted. For example, titanium is very noble and demonstrates excellent resistance to seawater, yet when less noble metals are coupled to titanium, galvanic corrosion usually is accelerated much less than would be anticipated, if at all. The reason for this is that titanium polarizes readily and quickly in seawater, thereby significantly reducing corrosion rate.
Factors Influencing Galvanic Corrosion Behavior Factors such as anode-to-cathode area ratios, distance between electrically connected materials, and geometric shapes also affect galvaniccorrosion behavior. Area effects in galvanic corrosion involve the ratio of the surface area of the more noble to the more active member(s). When the surface area of the more noble metal or alloy is large in comparison to the more active member, an unfavorable area ratio exists for the prevailing situation in which a couple is under cathodic control. The anodic current density on the more active metal or alloy is extremely large; therefore, the resulting polarization leads to more pronounced galvanic corrosion. The opposite area ratio—large active member surface, smaller noble member surface—produces only slightly accelerated galvanic effects because of the predominant polarization of the more noble material. Effect of Distance. Dissimilar metals in a galvanic couple that are in close physical proximity usually suffer greater galvanic effects than those that are farther apart. The distance effect is dependent on solution conductivity because the path of current flow is the primary consideration. Thus, if dissimilar pipes are butt welded with the electrolyte flowing
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Corrosion: Understanding the Basics
through them, the most severe corrosion will occur adjacent to the weld on the anodic member. Effect of Geometry. The geometry of the circuit also enters into the effect to the extent that current will not readily flow around corners. This is simply an extension of the principle described previously, in which the current takes the path of least resistance.
Situations That Promote Galvanic Attack Galvanic corrosion of the anodic member(s) of a couple may take the form of either general or localized corrosion, depending on the configuration of the couple, the nature of the films induced, and the nature of the metals or alloys involved. Dissimilar metals commonly are combined in engineering designs by mechanical or other means, for example, in heating or cooling coils in vessels, heat exchangers, or machinery. Such combinations often lead to galvanic corrosion. Examples are shown in Fig. 25 to 27. Nonmetallic Conductors. Less frequently recognized is the influence of nonmetallic conductors as cathodes in galvanic couples. Carbon brick in vessels is strongly cathodic to the common structural metals and alloys. Impervious graphite, especially in heat-exchanger applications, is cathodic to the less noble metals and alloys. Carbon-filled polymers can act as noble metals in a galvanic couple. Graphite-epoxy structures in aerospace applications must be adequately insulated from aluminum to prevent galvanic corrosion. Another example is the behavior of conductive films, such as mill scale (magnetite, Fe3O4) or iron sulfides on steel, or of lead sulfate on lead. Such films can be cathodic to the base metal exposed at breaks or pores in the scale or even to such extraneous items as valves or pumps in
Fig. 25
Galvanic corrosion of painted steel auto body panel in contact with stainless steel wheel opening molding
Forms of Corrosion: Recognition and Prevention
a piping system. As described in the following example, passive surface films can contribute to galvanic corrosion problems. Example 2: Galvanic Corrosion Occurring at the Same-Metal Couple (Ref 7). A well water copper piping system failed because of pitting corrosion. The water had a pH of 7.0 to 7.7, and pitting was caused by dissolved carbon dioxide. During failure analysis, researchers noticed that corrosion attack also occurred where a new replacement
Fig. 26
Fig. 27
Galvanic corrosion of steel pipe at brass fitting in humid marine atmosphere
Failure of the aluminum inner ring of an extruder (for plastics) cooling system due to galvanic corrosion. Note the severe deterioration adjacent to nozzle holes where brass nozzles had been inserted. Source: Nalco Chemical Company
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Corrosion: Understanding the Basics
copper pipe section was joined into a copper pipe section that was in service for several years. The two pieces of copper were of the same designation, C11000, which is electrolytic tough pitch (ETP) copper (99.95Cu-0.04O). The replacement pipe was only in service for about 2 months before significant galvanic corrosion occurred. The old section had a passive film and a noble electrochemical potential. The new section had an active potential surface not yet protected by an oxide film. Corrective Measures. This galvanic corrosion can be corrected by the following method: · Thoroughly cleaning the system so the old metal has the same potential as the new section · Electrically isolating the new section from the old section with a rubber tube and securing the tube to the metal piping with hose clamps until the new section has developed a passive film. After 2 months, the new metal can be soldered to the old metal with insignificant galvanic corrosion occurring. · Injecting a chemical (inhibitor) into the water system to accelerate the formation of a passive film on the new metal · Selecting an alternative material of construction for replacement piping that would be noble to the old metal piping. A candidate would be 90-10 copper-nickel alloy (C70600).
Laboratory Tests. Laboratory electrochemical polarization studies were carried out on a piece of new copper pipe, the old copper pipe, and a 90Cu-10Ni sample. The testing indicated that the electrochemical potential for the new copper specimen was –0.137 V versus a saturated calomel reference electrode. The 2-year-old scaled copper specimen was –0.070 V, and the new 90Cu-10Ni specimen was –0.061 V. Final Recommendations. The length of the replacement pipe would be approximately 6 ft (1.8 m), while the old piping would be several hundred feet. Although galvanically more noble, the surface area of the cathode (90Cu-10Ni) would be sufficiently small compared to the surface area of the anode (old copper piping) so that galvanic corrosion would be minimized. Metallic Coatings. Two types of metallic coatings are used in engineering design: noble metal coatings and sacrificial metal coatings. Noble metal coatings are used as barrier coatings over a more reactive metal. Galvanic corrosion of the substrate can occur at pores, damage sites, and edges in the noble metal coating. Sacrificial metal coatings provide cathodic protection of the more noble base metal, as in the case of galvanized steel or alclad aluminum. Cathodic Protection. Magnesium, zinc, and aluminum galvanic (sacrificial) anodes are used in a wide range of cathodic protection applications. The galvanic couple of the more active metal and a more noble
Forms of Corrosion: Recognition and Prevention
structure (usually steel, but sometimes aluminum, as in underground piping) provides galvanic (cathodic) protection, while accelerated corrosion of the sacrificial metal (anode) occurs. Chapter 10 contains information on the principles and applications of this method of corrosion prevention and the selection of anode materials. Metal Ion Deposition. Ions of a more noble metal may be reduced on the surface of a more active metal—for example, copper on aluminum or steel, or silver on copper. This process is also known as cementation, especially with regard to aluminum alloys. The resulting metallic deposit provides cathodic sites for further galvanic corrosion of the more active metal.
Prevention of Galvanic Corrosion A number of procedures or practices can be used to combat or minimize galvanic corrosion. Sometimes a single practice is sufficient, but a combination of several of the following may be required: · Select combinations of metals as close together as possible in the galvanic series suitable for the particular application or service environment. · Avoid combinations in which the area of the less noble material is relatively small. It is good practice to use the more noble metals for fasteners or other parts in equipment built largely of less corrosionresistant materials if dissimilar metals must be used. · Insulate dissimilar metals wherever practicable; if possible, insulate them completely. A common error in this regard concerns bolted joints such as two flanges, for example, a pipe to a valve, where the pipe might be steel or lead and the valve a different material. Polymeric washers under the bolt heads and nuts are assumed to insulate the two parts, yet the shank of the bolt touches the flanges. This problem is solved by placing plastic tubes over the bolt shanks, plus the washers, so that the bolts are isolated completely from the flanges. Figure 28 shows proper insulation for a bolted joint. If complete insulation cannot be achieved, a material such as paint or a plastic coating at joints (to increase the resistance of the circuit) will help. · Apply coatings with caution. When painting, for example, do not paint the less noble material without also coating the more noble material; otherwise, greatly accelerated attack may be concentrated at imperfections in coatings on the less noble metal. Keep such coatings in good repair. If only one surface can be painted, the more noble surface should be chosen to reduce or eliminate the cathode area. · In cases where the metals cannot be painted and are connected by a conductor external to the liquid, the electrical resistance of the liquid
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Corrosion: Understanding the Basics
Fig. 28
Proper insulation of a bolted joint
path may be increased by keeping the metals as far apart as possible. This is not practical in most cases. · If practical, add chemical inhibitors to the corrosion solution according to the nature of the solution to be inhibited. This reduces the corrosiveness of the environment. · Avoid joining materials well apart in the galvanic series by threaded connections because the threads will probably deteriorate excessively. As shown in Fig. 28, much of the effective wall thickness of the metal is cut away during the threading operation. In addition, spilled liquid or condensed moisture can collect and remain in the thread grooves. Brazed joints are preferred, using a brazing alloy more noble than at least one of the metals to be joined. Welded joints using welds of the same alloy are even better. · Employ cathodic protection measures. Magnesium, zinc, and aluminum galvanic (sacrificial) anodes are used in a wide range of cathodic protection applications. The galvanic couple of the more active metals and a more noble structure provides galvanic (cathodic) protection, while accelerated corrosion of the sacrificial metal (anode) occurs. Galvanized steel is composed of a thin layer of zinc on a steel substrate. The active zinc provides cathodic protection to exposed steel surfaces through beneficial galvanic action.
Erosion-Corrosion General Description Erosion-corrosion is the acceleration or increase in the rate of deterioration or attack on a metal because of mechanical wear or abrasive contributions in combination with corrosion. The combination of wear or
Forms of Corrosion: Recognition and Prevention
abrasion and corrosion results in more severe attack than would be realized with either mechanical or chemical corrosive action alone. Metal is removed from the surface as dissolved ions, as particles of solid corrosion products, or as elemental metal. The spectrum of erosion-corrosion ranges from primarily erosive attack, such as sandblasting, filing, or grinding of a metal surface, to primarily corrosion failures where the contribution of mechanical action is quite small. This section focuses on those instances where there is a fairly well-defined contribution from both mechanical and corrosive factors. Erosion-corrosion resulting from the relative movement between a corrosive fluid and the metal surface is discussed first, followed by a discussion of two special forms of erosion-corrosion, namely, cavitation and fretting corrosion. Erosion-corrosion is characterized in appearance by grooves, waves, rounded holes, and/or horseshoe-shaped grooves. Examples are shown in Fig. 29 and 30. Analysis of these marks can help determine the direction of flow. Affected areas are usually free of deposits and corrosion products, although corrosion products can sometimes be found if erosioncorrosion occurs intermittently, and/or the liquid flow rate is relatively low. Most metals are susceptible to erosion-corrosion under specific conditions. Metals that depend on a relatively thick protective coating of corrosion product for corrosion resistance are frequently subject to erosion-corrosion. This is due to the poor adhesion of these coatings relative to the thin films formed by the classical passive metals, such as stainless steels and titanium. Both stainless steels and titanium are relatively immune to erosion-corrosion in many environments. Metals that are soft and readily damaged or worn mechanically, such as copper and lead, are quite susceptible to erosion-corrosion. Even the noble or precious metals, such as silver, gold, and platinum, are subject to erosioncorrosion. Figure 31 shows a schematic of erosion-corrosion of a
Fig. 29
Erosion-corrosion of a cast stainless steel pump impeller after exposure to hot concentrated sulfuric acid with some solids present. Note the grooves, gullies, waves, and valleys common to erosion-corrosion damage.
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condenser tube wall. The direction of flow and the resulting attack where the protective film on the tube has broken down are indicated. All types of corrosive media generally can cause erosion-corrosion, including gases, aqueous solutions, organic systems, and liquid metals. For example, hot gases may oxidize a metal and then a high velocity blow off an otherwise protective scale. Solids in suspension in liquids (slurries) are particularly destructive from the standpoint of erosioncorrosion. Virtually anything that is exposed to a moving liquid is subject to erosioncorrosion. Examples include piping systems, particularly at bends, elbows, or wherever there is a change in flow direction or increase in turbulence; pumps; valves; centrifuges; tubular heat exchangers; impellers; and turbine blades.
Fig. 30
Horseshoe-shaped depressions on the internal surface of a brass heat exchanger tube caused by erosion-corrosion. Source: Nalco Chemical Company
Fig. 31
Schematic of erosion-corrosion of a condenser tube
Forms of Corrosion: Recognition and Prevention
Critical Factors Influencing Erosion-Corrosion Erosion-corrosion is a fairly complex failure mode influenced by both metal characteristics and environmental factors. Although some of these factors are interrelated, they are discussed separately insofar as possible. Surface Films. The nature and properties of the protective “films” that form on some metals and alloys are very important from the standpoint of resistance to erosion-corrosion. The ability of these films to protect the metal depends on the speed or ease with which they form when originally exposed to the environment, their resistance to mechanical damage or wear, and their rate of reformation when destroyed or damaged. A hard, dense, adherent, and continuous film provides better protection than one that is easily removed by mechanical means or that “wears off.” A brittle film that cracks or spalls under stress is not protective. The nature of the protective film that forms on a given metal depends on the specific environment to which it is exposed; this in turn, determines its resistance to erosion-corrosion by that fluid. Stainless steels depend heavily on a protective film (passivity) for their good resistance to corrosion. Consequently, these materials are vulnerable to erosion-corrosion. Figure 32 shows rapid attack due to erosion-corrosion of type 316 (18Cr-12Ni) stainless steel by a sulfuric acid/ferrous sulfate slurry moving at high velocity. The rate of deterioration is about 110 mm/year 4500 mils/year) at 55 °C (130 °F). This material showed no weight loss and was completely passive under stagnant conditions, as shown by the data point on the abscissa at approximately 60 °C (140 °F).
Fig. 32
Effect of temperature and copper ion addition on erosion-corrosion of type 316 stainless steel. The velocity of the sulfuric acid slurry was 12 m/s (39 ft/s)
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Lead depends on the formation of a lead sulfate/lead oxide protective surface film for long life in sulfuric acid environments; in many cases, more than 20 years of service may be obtained. Lead gains weight when exposed to sulfuric acid because of the surface coating or corrosion product formed (except in strong acid wherein the lead sulfate is soluble and not protective). However, lead valves failed in less than 1 week, and lead bends were rapidly attacked in a plant handling a 3% sulfuric acid solution at 90 °C (194 °F). As a result of these failures, erosion-corrosion tests were made; the results are plotted in Fig. 33. Under static conditions, the lead exhibited no deterioration (slight gain in weight), as shown by the points on the abscissa. Under high-velocity conditions, erosion-corrosion attack increased with temperature, as shown by the curve. Although rapidly attacked in solutions with a concentration of about 93% or higher, where lead sulfate is soluble, the protectiveness of the coating on lead can vary in dilute concentrations. Figure 34 shows that erosion-corrosion increases up to about 25% acid and then decreases.
Fig. 33
Fig. 34
Effect of temperature and velocity on attack of lead. The velocity of the 10% sulfuric acid was 12 m/s (39 ft/s)
Erosion-corrosion of lead as a function of sulfuric acid concentration. Velocity, 12 m/s (39 ft/s); temperature, 95 °C (203 °F)
Forms of Corrosion: Recognition and Prevention
Apparently, the rate of formation of the lead sulfate coating and/or its stability is not sufficient to decrease attack until concentrations greater than 25% are reached. The nature of the coating undoubtedly changes with acid concentration. Contaminants in sulfuric acid may result in soft and loose sulfates on lead, as is discussed later. Variations in the amount of attack on steel by water with different pH values and at different velocities can be attributed to the nature and composition of the surface scales formed. The scale on specimens exhibiting high rates of deterioration is porous, has poor adhesion, an is not protective. Below pH 5, the corrosion product film is increasingly more soluble because the pH is reduced (more acidic), and high rates of attack are observed as the protective oxide scale breaks down. In regions of low attack (higher pH and lower velocity), the scale formed on the steel is nonporous and has strong adhesion. This protective barrier film is stable and results in low corrosion rates. The behavior of steel and low-alloy steel tubes handling oils at high temperatures in petroleum refineries depends somewhat on the sulfide films formed. When the film “erodes away,” erosion-corrosion and rapid attack occur. For example, a normally tenacious sulfide film becomes porous and nonprotective when cyanides are present in these organic systems. Tests on copper and brass in sodium chloride solutions with and without oxygen showed that copper was attacked more than brass in oxygensaturated solutions. The copper was covered with a black and yellowbrown film, copper chloride (CuCl2). The brass was covered with a dark gray film, cupric oxide (CuO). The better resistance of the brass to attack was attributed to the greater stability or protectiveness of the dark gray film. Difficulty was encountered in obtaining reproducible results until a controlled alkali cleaning and drying procedure for the specimens was adopted. This indicates that surface films formed on copper and brass, because of atmospheric exposure, abrading, or other reasons, can have a definite effect on erosion-corrosion performance under some conditions. Titanium is a reactive metal but is resistant to erosion-corrosion in many environments because of the stability of the titanium dioxide (TiO2) film formed. It shows excellent resistance to seawater and chloride solutions and also to nitric acid. This corrosion resistance can be destroyed by erosion and wear. Strongly reducing conditions, such as deaerated hydrochloric acid, are damaging to titanium. Velocity. Because erosion-corrosion involves movement between a metal and its environment, the velocity of the environment plays an important role. Velocity often strongly influences the mechanism of the corrosion reaction. High velocities and environments containing solids in suspension result in a corrosion mechanism that is more mechanical in nature. Figures 32 and 33 show large increases in attack because of
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velocity. Table 3 shows the effect of velocity on a variety of metals and alloys exposed to seawater. These data show that the effect of velocity may be nil or extremely great. Increases in velocity generally result in increased attack, particularly if substantial rates of flow are involved. The effect may be nil or increase slowly until a critical velocity is reached, and then the attack may increase at a rapid rate. Figure 35 shows the “breakaway” or critical velocity effect. Table 3 lists several examples that exhibit little effect when the velocity is increased from 0.3 to 1.2 m/s (1 to 4 ft/s), but which undergo destructive attack at 8.2 m/s (27 ft/s). This high velocity is below the critical value for other materials listed at the bottom of Table 3. Erosion-corrosion can occur on metals and alloys that are completely resistant to a particular environment at low velocities. For example, hardened straight-chromium stainless steel valve seats and plugs give excellent service in most steam applications, but grooving or so-called “wire drawing” occurs in high-pressure steam reducing or throttling valves. Table 3 Erosion-corrosion rates of metals by seawater moving at different velocities Material
0.3 m/s (1 ft/s)(a)
Carbon steel Cast iron Silicon bronze Admiralty brass Hydraulic bronze G bronze Aluminum bronze (10% Al) Aluminum brass 90-10 Cu-Ni (0.8% Fe) 70-30 Cu-Ni (0.5% Fe) 70-30 Cu-Ni (0.5% Fe) Monel Stainless steel (Type 316) Hastelloy C Titanium
34 45 1 2 4 7 5 2 5 2 <1 <1 1 <1 0
Typical corrosion rates, mg/dm2/day 1.2 m/s (4 ft/s)(b) 8.2 m/s (27 ft/s)(c)
72 ¼ 2 20 1 2 ¼ ¼ ¼ ¼ <1 <1 0 ¼ ¼
254 270 343 170 339 280 236 199 99 199 39 4 <1 3 0
(a) Immersed in tidal current at a velocity of 0.3 m/s (1 ft/s). (b) Immersed in seawater flume at a velocity of 1.2 m/s (4 ft/s). (c) Attached to immersed rotating disk at a velocity of 8.2 m/s (27 ft/s)
Fig. 35
Schematic of the critical velocity effect for erosion-corrosion
Forms of Corrosion: Recognition and Prevention
Increased velocity may increase or reduce attack, depending on its effect of the corrosion mechanism involved. It may increase attack on steel by increasing the supply of oxygen, carbon dioxide, or hydrogen sulfide in contact with the metal surface. Velocity may also increase diffusion or transfer of ions by reducing the thickness of the stagnant film at the surface. Velocity can decrease attack and increase the effectiveness of inhibitors by supplying the chemical to the metal surface at a higher rate. It has been shown that less sodium nitrite inhibitor was needed at high velocity to protect steel in tap water. Similar mechanisms have been postulated for other types of inhibitors. It should be noted that the effective use of inhibitors to decrease erosion-corrosion depends, in many cases, on the nature and type of films formed on the metal as a result of the reaction between the metal and the inhibitor. Higher velocities may also decrease attack in some cases by preventing the deposition of silt or dirt that would cause concentration-cell corrosion. On the other hand, solids in suspension moving at high velocity may have a scouring effect and thus destroy surface protection. Many stainless steels have strong tendency to pit in seawater and other chlorides. However, some of these materials are used successfully in seawater, provided that the water is kept moving at a substantial velocity. This motion prevents the formation of deposits and concentration cells that cause pitting. Impingement. Impingement attack has been defined as localized erosion-corrosion caused by turbulence or impinging flow. This type of corrosion occurs in pumps, valves, orifices, on heat exchanger tubes, and at elbows and tees in tubes or pipelines. It occurs as deep, clean, horseshoe-shaped pits with the deep, or undercut, end pointing in the direction of flow. Impingement-corrosion attack can also occur as the result of partial blockage of a tube. A stone, a piece of wood, or some other object can cause the main flow to deflect against the wall of the tube, thereby creating turbulence and increasing fluid velocity. The impinging stream can rapidly perforate tube walls. Water that carries sand, silt, or mud will have an additional severely erosive effect on tubes. Steam erosion is another form of impingement corrosion. It occurs when high-velocity wet steam contacts a metal surface. The resulting attack usually produces a roughened surface showing a large number of small cones with the points facing in the direction of flow. Influence of Galvanic Corrosion. Galvanic, or two-metal corrosion, corrosion can influence erosion-corrosion when dissimilar metals are in contact in a flowing system. The galvanic effect may be nil under static conditions but may be greatly increased when movement is present. Figure 36 shows that attack on type 316 by itself was nil in highvelocity sulfuric acid, but increased to very high values when this alloy was in contact with lead. The passive film was destroyed by the
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combined effects of galvanic corrosion and erosion-corrosion. Couples of lead and type 316 showed no corrosion under static conditions. Velocity changes can produce surprising galvanic effects. In seawater at low velocity, the corrosion of steel is not appreciably affected by coupling with stainless steel, copper, nickel, or titanium. At high velocities, the attack on steel is much less when coupled to stainless steel and titanium than when coupled to copper or nickel. This is attributed to the more effective cathodic polarization of stainless steel and titanium at high velocities. Nature of Metal or Alloy. The chemical composition, corrosion resistance, hardness, and metallurgical history of metals and alloys can play important roles in the performance of these materials under erosioncorrosion conditions. The composition of most alloys determines to some extent their general corrosion resistance. If a reactive metal or an alloy composed of active elements is under consideration, its corrosion resistance is due chiefly to its ability to form and maintain a protective film—for example, aluminum and the stainless steels. The more noble metals possess good inherent corrosion resistance. In general, a material with better inherent resistance would be expected to show better performance when all other factors are equal. For example, an 80Ni-20Cr alloy should be superior to an 80Fe-20Cr alloy because nickel has better inherent corrosion resistance than iron. For the same reason, a nickel-copper alloy should be (and is) better than an alloy of zinc and copper. The addition of a third (or fourth) element to an alloy often increases its resistance to erosion-corrosion. For example, the addition of iron to a copper-nickel alloy produces a marked increase in resistance to erosion-
Fig. 36
Effect of contact with lead on erosion-corrosion of type 316 stainless steel. The velocity of the 10% sulfuric acid solution was 12 m/s (39 ft/s)
Forms of Corrosion: Recognition and Prevention
corrosion by seawater, as shown in Table 3. The addition of molybdenum to 18Cr-8Ni stainless steel (type 304) to make type 316 (18Cr-12Ni-2Mo) makes it more resistant to corrosion and erosioncorrosion. In both of these cases, the additional element produces a more stable protective film. Aluminum brasses show better erosioncorrosion resistance than straight (copper-zinc) brass. Reports on the resistance of steel and iron-chromium alloys to acid mine waters under erosion-corrosion conditions showed a straight-line increase in resistance with increasing chromium up to 13%. At this content and above, no attack occurred. Low-alloy chromium steels showed better erosion-corrosion resistance than straight carbon steels in high-temperature boiler feedwater. Chromium-plated steel and a leaded bronze were not suitable. Type 3 Ni-Resist, a cast iron containing 30% nickel (Ni) and 3% chromium (Cr), showed practically no attack by seawater after 60 days under erosion-corrosion conditions, whereas ordinary cast iron was badly deteriorated. Table 3 shows the effect of various material compositions on erosion-corrosion by seawater. Soft metals such as lead are, in a general way, more susceptible than the hard metals to erosion-corrosion because they are more subject to mechanical wear. Hardness is a fairly good criterion for predicting resistance to straight erosion or abrasion but it is not necessarily a good criterion for predicting resistance to erosion-corrosion. There are many methods for hardening metals and alloys. Perhaps the only method that comes close to producing good erosion-corrosion resistance is solidsolution hardening. This method involves adding an element to a given metal to produce a solid solution that is corrosion resistant and is inherently harder than the original metal. It cannot be softened by heat treatment or further hardened by heat treatment. Heat Treatment Effects. Hardening by heat treatment results in changes in microstructure and usually in heterogeneity. Quite often this decreases corrosion resistance because local anodic and cathodic areas are more readily formed. This makes a material more susceptible to erosion-corrosion. For example, the precipitation-hardening stainless steels would not be expected to perform as well as type 304 stainless steel under erosion-corrosion conditions, because the latter forms a more stable protective (passive) film. An example of poor performance by a high-hardness material is presented in the comparison of type 316 and type 329 stainless steels, the latter of which is hardenable by heat treatment. Both types of stainless steels show no measurable corrosion in a sulfuric acid slurry under static conditions, even when the type 329 is age hardened to 450 HB (Brinell hardness scale). Under the erosion-corrosion conditions in a centrifuge test, however, the hard type 329 steel deteriorates more than 10 times faster than the soft (150 HB) type 316. Both materials possess good “passive” resistance in the static condition, but the type 316 steel has a more stable
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film and maintains passivity more successfully when exposed to the more severe erosion-corrosion conditions in the centrifuge test. The nature of the carbide structure in carbon steel may have an important bearing on its resistance to erosion-corrosion. In some oil well and sulfuric acid applications involving flow, so-called “ringworm corrosion” occurs. Carbon steel with a normal pearlitic structure is not attacked significantly, but the same steel containing spheroidized pearlite is badly corroded. Cast iron sometimes shows better performance than steel under erosioncorrosion conditions, particularly in hot, strong sulfuric acid. The iron in the cast iron is corroded, but the remaining “graphitized” layer, consisting of the original graphite network and corrosion products, forms a protective layer.
Prevention of Erosion-Corrosion Five methods are used to prevent or minimize damage due to erosioncorrosion. In order of their extent of use, they are use of materials with better resistance to erosion-corrosion, design, alteration of the environment, coatings, and cathodic protection. The reasons for using better materials that give better performance are obvious. This method usually represents an economical solution to most erosion-corrosion problems. Design is an important corrosion-control method in that the life of presently used or less costly materials can be extended considerably or the attack practically eliminated. Design here involves change in shape or geometry, and not in selection of material. Erosion-corrosion damage can be reduced by implementing the following design recommendations: · Increase the pipe diameter to decrease velocity not only from the mechanical standpoint but also to ensure lamellar flow. · Increase the diameter and/or streamline bends to reduce impingement effects. · Increase the thickness of material in vulnerable areas. In one instance of severe erosion-corrosion of lead, maintenance costs were reduced to a satisfactory level by using a sweeping bend and doubling the thickness of the pipe. · Streamline the design of other equipment, such as inlets and outlets, to remove obstructions for the same reasons. · Insert readily replaceable impingement plates or baffles. · Direct inlet pipes toward the center of a tank instead of near its wall. · Design tubes to extend several centimeters beyond the tube sheet at the inlet end. In one case, tubing life was practically doubled by increasing the length by 10 cm (4 in.). The protruding tube ends were attacked, but operation was not affected.
Forms of Corrosion: Recognition and Prevention
· Insert ferrules or short lengths of flared tubing in the inlet ends. These could be made of the same material as the tubes or of materials with better resistance. Nonmetallic ferrules are available and are used in condensers. The end of the ferrule should be “feathered” to blend the flow. If this is not done, erosion-corrosion occurs on the tube just beyond the end of the ferrule because of the “step” present. Galvanic corrosion must be considered when using metallic inserts. Tubing life in a vertical evaporator was doubled by turning the evaporator upside down when the inlet or bottom ends of the tubes became thin. The outlet ends, which were not appreciably attacked, became inlet ends. · Design equipment so that parts are readily replaceable. For instance, tube bundles that are readily removed and replaced by spares can be repaired at leisure. Buckets and conveyor flights that are easily installed on centrifugals and other conveying equipment reduce costs. Use of pumps with interchangeable parts of different alloys helps to reduce costs when an unsatisfactory alloy is originally selected. When shutdown of equipment is serious from a production standpoint, it is often desirable to design a standby or alternate unit.
Good design implies proper construction and workmanship. In one case, some of the blades of a steam turbine were out of line, and the protruding blades suffered severe erosion-corrosion damage from water droplets in the stream. Misalignment from one pipe section to the next can cause erosion-corrosion in both flanged and welded joints. Environment. Alteration of the environment, such as deaeration and the addition of inhibitors, is effective and is sometimes used to minimize erosion-corrosion damage. In many cases, however, this method is not economical. Settling and/or filtration to remove solids is very helpful. Whenever possible, the temperature of the environment should be reduced. This has been done in many cases without appreciably affecting the particular process. Temperature is a serious problem in erosion- corrosion, as it is in all types of corrosion. Coatings. Applied coatings of various types that produce an effective barrier between the metal and its environment can be used, but this method is not always feasible for solving most erosion-corrosion problems. Hardfacings, or weld overlays, are sometimes helpful, provided that the facing has good corrosion resistance. Repair or buildup of attacked areas by welding is often feasible. Cathodic protection has been used to reduce erosion-corrosion attack, but it has not found widespread application. Cathodic protection systems for preventing erosion-corrosion must be properly designed and maintained to be effective. Corrosion can be intensified if the polarity of the cathodic protection system is inadvertently reversed.
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Cavitation General Description. Cavitation is a form of erosion-corrosion that is caused by the formation and collapse of vapor bubbles in a liquid against a metal surface. Cavitation occurs in hydraulic turbines, on pump impellers, on ship propellers, and on many surfaces in contact with high-velocity liquids subject to changes in pressure. Mechanism (Ref 8). Figure 37 is a simplified representation of the cavitation process. Figure 37(a) shows a vessel containing a liquid. The vessel is closed by an airtight plunger. When the plunger is withdrawn (Fig. 37b), a partial vacuum is created above the liquid, causing vapor bubbles to form and grow within the liquid. In essence, the liquid boils without a temperature increase. If the plunger is then driven toward the surface of the liquid (Fig. 37c), the pressure in the liquid increases, and the bubbles condense and collapse (implode). In a cavitating liquid, these three steps occur in a matter of milliseconds. As shown in Fig. 37(d), implosion of a vapor bubble creates a microscopic “torpedo” of water that is ejected from the collapsing bubble at velocities that may range from 100 to 500 m/s (330 to 1650 ft/s). When the torpedo impacts the metal surface, it dislodges protective surface films and/or locally deforms the metal itself. Thus, fresh surfaces are exposed to corrosion and the reformation of protective films, which is followed by more cavitation, and so on. Damage occurs when the cycle is allowed to repeat over and over again. Partial vacuum
Pressurized
Metal (a) Rest Quiescent liquid at standard temperature and pressure
(b) Expansion Liquid boiling at room temperature
(c) Compression Collapse of vapor bubbles
Metal oxide (d)
Approaching microjet torpedo
Fig. 37
Destruction of metal oxide on impact
Repair of metal oxide at expense of metal
Schematic representation of cavitation showing a cross section through a vessel and plunger enclosing a fluid. (a) Plunger stationary, liquid at standard temperature and pressure. (b) Plunger withdrawn, liquid boils at room temperature. (c) Plunger advanced, bubbles collapse. (d) Disintegration of protective corrosion product by impacting microjet “torpedo.” Source: Ref 8
Forms of Corrosion: Recognition and Prevention
The appearance of cavitation is similar to pitting except that surfaces in the pits are usually much rougher. The affected region is free of deposits and accumulated corrosion products if cavitation has been recent. Typically, affected surfaces are highly localized to specific regions, although, if cavitation is severe and widespread, the area affected may be extensive (Fig. 38). On symmetrical components having repeated elements (e.g., impellers) the pattern of damage may repeat itself at identical locations on each element as shown in Fig. 39. Prevention. Cavitation can be controlled or minimized by the following: · Improving design to minimize hydrodynamic pressure differences · Employing stronger (harder) and more corrosion-resistant materials. Figure 40 compares the cavitation resistance of a number of metals and alloys. Cobalt-based Stellite alloys are frequently employed for applications requiring resistance to cavitation. Cobalt-containing stainless steels have also proved effective. Hardfacing alloys are also frequently employed. · Specifying a smooth finish on all critical metal surfaces · Coating with resilient materials such as rubber or some plastics. However, such materials are subject to disbondment at the metal/ elastomer interface at high cavitation intensities, even if the exposure is brief.
Fig. 38
A cast steel feedwater-pump impeller severely damaged by cavitation. Note how damage is confined to the outer edges of the impeller where vane speed was maximum. Source: Nalco Chemical Company
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Fig. 39
Fig. 40
Cavitation damage repeated on successive vanes of a bronze impeller. Source: Nalco Chemical Company
Classification of 22 alloys or alloy groups according to their normalized cavitation erosion resistances relative to 18Cr-8Ni austenitic stainless steel having a hardness of 170 HV
Forms of Corrosion: Recognition and Prevention
Fretting Corrosion General Description. Fretting corrosion is a combined wear and corrosion process in which material is removed from contacting surfaces when motion between the surfaces is restricted to very small amplitude oscillations (often, the relative movement is barely discernible). Usually, the condition exists in machine components that are considered fixed and not expected to wear. Pressed-on wheels can often fret at the shaft/wheel hole interface. Oxidation is the most common element in the fretting process. In oxidizing systems, fine metal particles removed by adhesive wear are oxidized and trapped between the fretting surfaces (Fig. 41). The oxides act like an abrasive (such as lapping rouge) and increase the rate of material removal. This type of fretting in ferrous alloys is easily recognized by the red material oozing from between the contacting surfaces. Fretting damage can be very serious and catastrophic. For example, in gas turbines, the blades are often anchored in dovetail slots in the hub. Vibrations in the blade system can cause fretting in the blade roots and loosening. The loose blades then rattle around and break out because of fatigue or impact wear. A piece of blade moving into the gas flow through the turbine can cause severe damage, destroying rows of blades and causing engine failure. Another fretting problem that has injury potential is the loosening of wheels or flywheels from shafts or axles. Railroad car wheels, for example, are shrink fitted onto their axles. If the wheel loosens from running vibrations and comes off during operation of the train, it can cause derailment and has the potential to become a loose rolling missile capable of penetrating nearby buildings. Other mechanical parts susceptible to fretting damage include couplings, riveted and pin joints, surgical implants, rolling contact bearings, and bridge bearings. Recognizing Fretting Corrosion. Fretting corrosion takes the form of local surface dislocations and deep pits. These occur in regions where slight relative movements have occurred between mating, highly loaded surfaces. The pits can sometimes provide stress raisers for the
Fig. 41
Schematic of the fretting process.
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Fig. 42
Fretting corrosion of the bearing race of a helicopter drive train over-running clutch. This problem was caused by vibration (and rubbing) of the ball in the inner and outer races of bearings that support the rotor shaft. Note the two areas on the left- and right-hand sides of the figure exhibiting fretting damage.
Fig. 43
Fretting corrosion on the root surface of an aircraft power plant compressor blade. Fatigue cracks can initiate as a result of this fretting pitting damage.
initiation of corrosion fatigue. Examples of fretting corrosion of aircraft components are shown in Fig. 42 and 43. Prevention. Fretting corrosion can be controlled by the following · Lubricating (e.g., low-viscosity oils) the faying surfaces · Restricting the degree of movement · Selecting materials and combinations that are less susceptible to fretting corrosion (Table 4)
Forms of Corrosion: Recognition and Prevention
151
Table 4 Resistance to fretting corrosion of various material couples under dry conditions Couple
Steel on steel Nickel on steel Aluminum on steel Antimony plate on steel Tin on steel Aluminum on aluminum Zinc-plated steel on aluminum Iron-plated steel on aluminum Cadmium on steel Zinc on steel Copper alloys on steel Zinc on aluminum Copper plate on aluminum Nickel plate on aluminum Iron plate on aluminum Silver plate on aluminum Lead on steel Silver plate on steel Silver plate on aluminum plate Steel with conversion coating on steel
Resistance to fretting
Low Low Low Low Low Low Low Low Medium Medium Medium Medium Medium Medium Medium Medium High High High High
Intergranular Corrosion General Description Intergranular corrosion is defined as the selective dissolution of grain boundaries, or closely adjacent regions, without appreciable attack of the grains themselves. This dissolution is caused by potential differences between the grain-boundary region and any precipitates, intermetallic phases, or impurities that form at the grain boundaries. The actual mechanism differs with each alloy system. Precipitates that form as a result of the exposure of metals at elevated temperatures (for example, during production, fabrication, and welding) often nucleate and grow preferentially at grain boundaries. If these precipitates are rich in alloying elements that are essential for corrosion resistance, the regions adjacent to the grain boundary are depleted of these elements. The metal is thus sensitized and is susceptible to intergranular attack in a corrosive environment. For example, in austenitic stainless steels such as AISI type 304, the cause of intergranular attack is the precipitation of chromium-rich carbides ((Cr, Fe)23C6) at grain boundaries. These chromium-rich precipitates are surrounded by metal that is depleted in chromium; therefore, they are more rapidly attacked at these zones than on undepleted metal surfaces. Impurities that segregate at grain boundaries may promote galvanic action in a corrosive environment by serving as anodic or cathodic sites. Therefore, this would affect the rate of the dissolution of the alloy matrix in the vicinity of the grain boundary. An example of this is found in aluminum alloys that contain intermetallic compounds, such as Mg5Al8
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and CuAl2, at the grain boundaries. During exposures to chloride solutions, the galvanic couples formed between these precipitates and the alloy matrix can lead to severe intergranular attack. Susceptibility to intergranular attack depends on the corrosive solution and on the extent of intergranular precipitation, which is a function of alloy composition, fabrication, and heat treatment parameters.
Intergranular Corrosion of Austenitic Stainless Steels General Description. At temperatures above about 1035 °C (1900 °F), chromium carbides are completely dissolved in austenitic stainless steels. However, when these steels are slowly cooled from these high temperatures or reheated into the range of 425 to 815 °C (800 to 1500 °F), chromium carbides are precipitated at the grain boundaries. These carbides contain more chromium than the matrix does. The precipitation of the carbides depletes the matrix of chromium adjacent to the grain boundary. The diffusion rate of chromium in austenite is slow at the precipitation temperatures; therefore, the depleted zone persists, and the alloy is sensitized to intergranular corrosion. This sensitization occurs because the depleted zones have higher corrosion rates than the matrix in many environments. Figure 44 illustrates how the chromium content influences the corrosion rate of iron-chromium alloys in boiling 50% sulfuric acid (H2SO4) containing ferric sulfate (Fe2(SO4)3). In all cases, the alloys are in the passive state. The wide differences in the corrosion rate are the result of the differences in the chromium content. If the austenitic stainless steels are cooled rapidly to below about 425 °C (800 °F), the carbides do not precipitate, and the steels are immune to intergranular corrosion. Reheating the alloys to 425 to 815 °C (800 to 1500 °F), as for stress relief, will cause carbide precipitation and sensitivity to intergranular corrosion. The maximum rate of carbide precipitation occurs at about 675 °C (1250 °F). Because this is a common temperature for the stress relief of carbon and low-alloy steels, care must be exercised in selecting stainless steels to be used in dissimilar-metal joints that are to be stress relieved. Effects of Welding. Welding is the common cause of the sensitization of stainless steels to intergranular corrosion. Although the cooling rates in the weld itself and the base metal immediately adjacent to it are sufficiently high to avoid carbide precipitation, the weld thermal cycle will bring part of the heat-affected zone (HAZ) into the precipitation range. Carbides will precipitate, and a zone somewhat removed from the weld will become susceptible to intergranular corrosion (Fig. 45). Welding does not always sensitize austenitic stainless steels. In thin sections, the thermal cycle may be such that no part of the HAZ is at sensitizing temperatures long enough to cause carbide precipitation. Once
Forms of Corrosion: Recognition and Prevention
Fig. 44
Fig. 45
The effect of chromium content on the corrosion behavior of ironchromium alloys in boiling 50% H2SO4 with Fe2(SO4)3.
Intergranular corrosion of sensitized HAZ grain boundaries and methods for its prevention. The four different panels were joined by welding and then exposed to a hot solution of nitric-hydrofluoric acid (HNO3-HF). Weld decay, such as that shown in the type 304 steel (bottom right), is prevented by reduction of the carbon content (type 304L, top left) or by stabilization with titanium (type 321, bottom left) or niobium (type 347, top right).
153
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Corrosion: Understanding the Basics
the precipitation has occurred, it can be removed by reheating the alloy to above 1035 °C (1895 °F) and cooling it rapidly. Prevention. Susceptibility to intergranular corrosion in austenitic stainless steels can be avoided by controlling their carbon contents or by adding elements whose carbides are more stable than those of chromium. For most austenitic stainless steels, restricting their carbon contents to 0.03% or less will prevent sensitization during welding and most heat treatment. This method is not effective for eliminating sensitization that would result from long-term service exposure at 425 to 815 °C (800 to 1500 °F). Titanium and niobium form more stable carbides than chromium and are added to stainless steels to form these stable carbides, which remove carbon from solid solution and prevent precipitation of chromium carbides. The most common of these stabilized grades are types 321 and 347. Type 321 contains a minimum of titanium equal to 5 ´ (C + N)%, and type 347, a minimum of niobium equal to 8 ´ C%. Nitrogen must be considered when titanium is used as a stabilizer, not because the precipitation of chromium nitride is a problem in austenitic steels, but because titanium nitride is very stable. Titanium will combine with any available nitrogen; therefore, this reaction must be considered when determining the total amount of titanium required to combine with the carbon. The stabilized grades are more resistant to sensitization by longterm exposure at 425 to 815 °C (800 to 1500 °F) than the low-carbon grades are, and the stabilized grades are the preferred materials when service involves exposure at these temperatures. For maximum resistance to intergranular corrosion, these grades are given a stabilizing heat treatment at about 900 °C (1650 °F). The purpose of treatment is to remove carbon from solution at temperatures where titanium carbides are stable but chromium carbides are not. Such treatments prevent the formation of chromium carbide when the steel is exposed to lower temperatures. Figure 45 illustrates how both carbon control and stabilization can eliminate intergranular corrosion in as-welded austenitic stainless steels. It also shows that the sensitized zone in these steels is somewhat removed from the weld metal. Knife-Line Attack. Stabilized austenitic stainless steels may become susceptible to a localized form of intergranular corrosion known as knife-line attack or knife-line corrosion. During welding, the base metal immediately adjacent to the fusion line is heated to temperatures high enough to dissolve the stabilizing carbides, but the cooling rate is rapid enough to prevent carbide precipitation. Subsequent welding passes reheat this narrow area into the temperature range in which both the stabilizing carbide and the chromium carbide can precipitate. The
Forms of Corrosion: Recognition and Prevention
precipitation of chromium carbide leaves the narrow band adjacent to the fusion line susceptible to intergranular corrosion. Knife-line attack can be avoided by the proper choice of welding variables and by the use of stabilizing heat treatments. Testing. The common methods of testing austenitic stainless steels for susceptibility to intergranular corrosion are described in ASTM A 262 “Standard Practices for Detecting Susceptibility to Intergranular Attack in Austenitic Steels.” Practice A is a screening test that uses an electrolytic oxalic acid etch combined with metallographic examination. The other practices (B through E) involve exposing the material (possibly after a sensitizing treatment) to boiling solutions of 65% nitric acid (HNO3), acidified ferric sulfate (Fe2(SO4)3) solution, nitrichydrofluoric acid (HNO3-HF) solution, or acidified cupric sulfate (CuSO4) solution, depending on the specific alloy and its application. Similar ASTM tests have been developed for other higher-alloyed austenitic stainless steels, ferritic stainless steels, high nickel-based alloys, and aluminum alloys (Table 5).
Intergranular Corrosion of Other Alloy Systems Ferritic Stainless Steels and Nickel-Base Alloys. Some ferritic stainless steels (e.g., types 430 and 446) are susceptible to intergranular corrosion, as are the following nickel alloys: Inconel alloys 600 and 601, Incoloy 800, despite the presence of titanium, Incoloy 800H, Nickel 200, and Hastelloy alloys B and C. Intergranular corrosion in these alloys is avoided by one or a combination of the following: · Keeping the alloy in the solution heat-treated condition at all times. · Limiting interstitial elements, primarily carbon and nitrogen, to the lowest practical levels. · Adding carbide-stabilizing elements, such as titanium, niobium, and titanium, along with a stabilizing heat treatment where necessary.
Aluminum Alloys. Intergranular corrosion in aluminum alloys is caused by potential differences between the grain-boundary region and the adjacent grain bodies. The location of the anodic path varies with the different alloy systems. In 2xxx series alloys, it is a narrow band on either side of the grain boundary that is depleted in copper; in 5xxx alloys, it is the anodic constituent MgAl3, when that constituent forms a continuous path along a grain boundary; in copper-free 7xxx series alloys, it is generally considered to be the anodic zinc- and magnesiumbearing constituents on the grain boundary; and in the copper-bearing 7xxx series alloys, it is the copper-depleted bands along the grain
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Corrosion: Understanding the Basics
boundaries. The 6xxx series alloys generally resist this type of corrosion, although slight intergranular attack has been observed in highly aggressive environments. The 1xxx and 3xxx series are immune to intergranular corrosion. Other alloy systems such as copper, magnesium, lead, and zinc are also susceptible to intergranular corrosion under very specific conditions. However, very few case histories have been reported in the literature.
Table 5
Appropriate evaluation tests and acceptance criteria for wrought alloys
Sensitizing treatment
Exposure time, h
Ferric sulfate (A 763-X) Ferric sulfate (A 763-X) Ferric sulfate (A 763-X)
None None None
24 72 120
None None None None
Type 321 Type 347 20Cb-3 904L Incoloy 825 Hastelloy G
Cupric sulfate (A 763-Y) Ferric sulfate (A 763-X) Ferric sulfate (A 763-X) Oxalic acid (A 262-A) Ferric sulfate (A 262-B) Oxalic acid (A 262-A) Nitric acid (A 262-C) Nitric acid (A 262-C) Oxalic acid (A 262-A) Ferric sulfate (A 262-B) Oxalic acid (A 262-A) Ferric sulfate (A 262-B) Oxalic acid (A 262-A) Ferric sulfate (A 262-B) Oxalic acid (A 262-A) Ferric sulfate (A 262-B) Nitric acid (A 262-C) Nitric acid (A 262-C) Ferric sulfate (G 28-A) Ferric sulfate (G 28-A) Nitric acid (A 262-C) Ferric sulfate (G 28-A)
1 h at 675 °C (1250 °F) 1 h at 675 °C (1250 °F) 1 h at 675 °C (1250 °F) None 1 h at 675 °C (1250 °F) None
120 120 120 ¼ 120 ¼ 240 240 ¼ 120 ¼ 120 ¼ 120 ¼ 120 240 240 120 120 240 120
N06985
Hastelloy G-3
Ferric sulfate (G 28-4)
None
120
N06625 N06690 N10276 N06455 N06110 N10001
Inconel 625 Inconel 690 Hastelloy C-276 Hastelloy C-4 Allcorr Hastelloy B
Ferric sulfate (G 28-A) Nitric acid (A 262-C) Ferric sulfate (G 28-A) Ferric sulfate (G 28-A) Ferric sulfate (G 28-B) 20% Hydrochloric acid
None 1 h at 540 °C (1000 °F) None None None None
120 240 24 24 24 24
N10665
Hastelloy B-2
20% Hydrochloric acid
None
24
Concentrated nitric acid (G 67)
None
24
1.14 (45) 0.25 (10) 0.05 (2) and no significant grain dropping No significant grain dropping No significant grain dropping No significant grain dropping (a) 0.1 (4) (a) 0.05 (2) 0.025 (1) (a) 0.1 (4) (a) 0.1 (4) (a) 0.1 (4) (a) 0.1 (4) 0.05 (2) 0.05 (2) 0.05 (2) 0.05 (2) 0.075 (3) 0.043 (1.7) sheet, plate, and bar; 0.05 (2) pipe and tubing 0.043 (1.7) sheet, plate, and bar; 0.05 (2) pipe and tubing 0.075 (3) 0.025 (1) 1 (40) 0.43 (17) 0.64 (25) 0.075 (3) sheet, plate, and bar; 0.1 (4) pipe and tubing 0.05 (2) sheet, plate, and bar; 0.086 (3.4) pipe and tubing (b)
Sodium chloride + hydrogen peroxide (G 110)
None
24
(c)
UNS number
Alloy name
S43000 S44600 S44625
Type 430 Type 446 26-1
S44626 S44700 S44800 S30400
26-1S 29-4 29-4-2 Type 304
S30403
Type 304L
S30908 S31600
Type 309S Type 316
S31603
Type 316L
S31700
Type 317
S31703
Type 317L
S32100 S34700 N08020 N08904 N08825 N06007
A95005-A95657
Aluminum Association 5xxx alloys A92001-A92618, Aluminum A97001-A97475 Association 2xxx and 7xxx alloys
Applicable tests (ASTM standards)
Criteria for passing, appearance, or maximum allowable corrosion rate, mm/month (mils/month)
(a) See A 262, practice A. (b) See G 67, section 4.1. (c) See G 110
1 h at 675 °C (1250 °F) None None 1 h at 675 °C (1250 °F) None 1 h at 675 °C (1250 °F)
Forms of Corrosion: Recognition and Prevention
Exfoliation General Description. Exfoliation is a form of localized corrosion that primarily affects aluminum alloys in industrial or marine environments. Corrosion proceeds laterally from initiation sites on the surface and generally proceeds intergranually along planes parallel to the surface. The corrosion products that form in the grain boundaries force metal away from the underlying base material, resulting in a layered or flakelike appearance (Fig. 46). Prevention Through Proper Alloy/Temper Selection. The commercialpurity (1xxx) and aluminum-manganese (3xxx) alloys are quite resistant to exfoliation corrosion in all tempers. Exfoliation has been encountered in some highly cold-worked aluminum-magnesium (5xxx) materials such as 5456-H321 boat hull plates. These developed a highly elongated grain structure and selective grain boundary precipitation. This exfoliation problem led to the establishment of special boat hull plate tempers, H116 and H117, for alloys 5083, 5086, and 5456, which have high resistance to exfoliation corrosion. In the heat treatable aluminum-copper-magnesium (2xxx) and aluminumzinc-magnesium-copper (7xxx) alloys, exfoliation corrosion has usually been confined to relatively thin sections of highly worked products with an elongated grain structure. In 2124-T351 plate, for example,
Fig. 46
Exfoliation corrosion in an alloy 7178-T651 plate exposed to a seacoast environment. Cross section of the plate shows how exfoliation develops by corrosion along boundaries of thin, elongated grains. As corrosion proceeds along multiple narrow paths parallel to the surface, the insoluble products that are formed occupy a larger volume than the metal consumed in producing them. These voluminous corrosion products exert a wedging action, which develops lateral tensile forces. This results in the splitting, flaking, or delamination of uncorroded layers of metal.
157
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Corrosion: Understanding the Basics
13 mm (0.5 in.) plate was quite susceptible in laboratory and atmospheric tests, while 50 mm (2 in.) and 100 mm (4 in.) plate, with less directional structures, did not exfoliate. In extrusions, the surface is often quite resistant to exfoliation because of its recrystallized grain structure. Subsurface grains are unrecrystallized, elongate, and vulnerable to exfoliation. In aluminum-zinc-magnesium alloys containing copper, such as 7075, resistance to exfoliation can be improved markedly by overaging. This is designated by the temper designations of T7xxx for wrought products. (e.g., 7075-T7351). While a 5 to 10% loss in strength occurs, improved resistance to exfoliation is provided. In copper-free or low-copper 7xxx alloys, exfoliation corrosion resistance can be controlled by overaging or by recrystallizing heat treatments and can also be controlled to some extent by changes in alloying elements. In aluminum-copper-magnesium (2xxx) alloys, artificial aging to the T6 or T8 condition provides improved resistance.
Dealloying Corrosion Dealloying, also referred to as selective leaching or parting corrosion, is a corrosion process in which the more active metal is selectively removed from an alloy, leaving behind a porous weak deposit of the more noble metal. Specific categories of dealloying often carry the name of the dissolved element. For example, the preferential leaching of zinc from brass is called dezincification. If aluminum is removed, the process is called dealuminification, and so forth. In the case of gray iron, dealloying is called graphitic corrosion. In the dealloying process, typically one of two mechanisms occurs: alloy dissolution and replating of the cathodic element or selective dissolution of an anodic alloy constituent. In either case, the metal is left spongy and porous and loses much of its strength, hardness, and ductility. Table 6 lists some of the alloy-environment combinations for which dealloying has been reported. Descriptions of the most common types of dealloying, dezincification and graphitic corrosion, follow.
Dezincification General Description. Copper-zinc alloys containing more than 15% zinc are susceptible to dezincification. In the dezincification of brass, selective removal of zinc leaves a relatively porous and weak layer of copper and copper oxide. Corrosion of a similar nature continues beneath the primary corrosion layer, resulting in gradual replacement of sound brass by weak, porous copper. Unless arrested, dealloying even-
Forms of Corrosion: Recognition and Prevention
159
tually penetrates the metal, weakening it structurally and allowing liquids or gases to leak through the porous mass in the remaining structure. Dezincification is the usual form of corrosion for uninhibited brasses in prolonged contact with waters high in oxygen and carbon dioxide (CO2). It is frequently encountered with quiescent or slowly moving solutions. Slightly acidic water, low in salt content and at room temperature, is likely to produce uniform attack, but neutral or alkaline water, high in salt content and above room temperature, often produces plug-type attack. Identification. The two major forms of dezincification are layertype and plug-type dealloying. In layer-type dezincification, the majority of the component surface is converted to corrosion product to a roughly uniform depth (Fig. 47 and 48). Plug-type dezincification produces small pockets or plugs of almost pure copper (Fig. 49 and 50). Layer-type dezincification is easy to recognize visually. The original component shape and dimensions are usually preserved but the metal changes from the gold yellow of zinc brass to the red of elemental copper at the dezincified surface (Fig. 47). Plug-type dezincification is common on horizontal pipe/tube sections. Attack is usually confined to areas beneath deposits lying along pipe bottoms (see Fig. 49). Plugs are sometimes blown out of the pipe wall due to internal pressure and produce holes. More often, fluids weep through the porous plugs, causing additional corrosion problems. Prevention through Alloy Selection. Brasses with copper contents of 85% or more resist dezincification. Dezincification of brasses with Table 6 Combinations of alloys and environments subject to dealloying and elements preferentially removed Alloy
Brasses
Environment
Copper-gold single crystals Monels
Many waters, especially under stagnant conditions Soils, many waters Hydrofluoric acid, acids containing chloride ions High-temperature steam and acidic species Hot brine or steam High heat flux and low water velocity (in refinery condenser tubes) Ferric chloride Hydrofluoric and other acids
Gold alloys with copper or silver Tungsten carbide-cobalt High-nickel alloys
Sulfide solutions, human saliva Deionized water Molten salts
Medium- and high-carbon steels
Oxidizing atmospheres, hydrogen at high temperatures High-temperature oxidizing atmospheres Oxygen at high temperature
Gray iron Aluminum bronzes Silicon bronzes Tin bronzes Copper nickels
Iron-chromium alloys Nickel-molybdenum alloys
Element removed
Zinc (dezincification) Iron (graphic corrosion) Aluminum (dealuminification) Silicon (desiliconification) Tin (destannification) Nickel (denickelification)
Copper Copper in some acids, and nickel in others Copper, silver Cobalt Chromium, iron, molybdenum, and tungsten Carbon (decarburization) Chromium, which forms a protective film Molybdenum
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Corrosion: Understanding the Basics
Fig. 47
Layer-type dezincification of a brass pump component. The dark (red) outer layers are uniformly corroded regions surrounding the uncorroded metal (original yellow) of the brass. Source: Nalco Chemical Company
Fig. 48
Layer-type dezincification of a thin brass sheet. The 0.48 mm (0.019 in.) sheet is shown in cross section. The dezincified layers converge toward the edge (left side) of the sheet. Note the porosity of the dezincified metal. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
two-phase structures is generally more severe, particularly if the second phase is continuous. It usually occurs in two stages: the high-zinc b phase, followed by the lower-zinc a phase. Tin tends to inhibit dealloying, especially in cast alloys. Alloys C46400 (naval brass) and C67500 (manganese bronze), which are a-b
Fig. 49
A large plug of dezincified metal beneath a deposit on a brass pipe. Source: Nalco Chemical Company
Plug-type dezincification in an a-brass (70Cu-30Zn) exposed for 79 days in 1 N NaCl at room temperature. Note porous structure within the plug. The dark line surrounding the plug is an etching artifact. 160´
Fig. 50
161
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Corrosion: Understanding the Basics
brasses containing about 1% Sn, are widely used for naval equipment and have reasonably good resistance to dezincification. Addition of a small amount of phosphorus, arsenic, or antimony to admiralty metal (an all-a 71Cu-28Zn-1Sn brass) inhibits dezincification. Inhibitors are not entirely effective in preventing dezincification of the a-b brasses because they do not prevent dezincification of the b phase. Where dezincification is a problem, red brass, commercial bronze, inhibited admiralty metal, and inhibited aluminum brass can be successfully used. In some cases, the economic penalty of avoiding dealloying by selecting a low-zinc alloy may be unacceptable. Low-zinc alloy tubing requires fittings that are available only as sand castings, but fittings for higher-zinc tube can be die cast or forged much more economically. Where selection of a low-zinc alloy is unacceptable, inhibited yellow brasses are generally preferred.
Graphitic Corrosion General Description. Graphitic corrosion is a form of dealloying unique to cast irons. It is observed in gray cast irons in relatively mild environments in which selective leaching of iron leaves a graphite network. Selective leaching of the iron takes place because the graphite is cathodic to iron, and the gray iron structure establishes an excellent galvanic cell. This form of corrosion generally occurs only when corrosion rates are low. If the metal corrodes more rapidly, the entire surface, including the graphite, is removed, and more or less uniform corrosion occurs. Graphitic corrosion can cause significant problems because, although no dimensional changes occur, the cast iron loses its strength and metallic properties. Thus, without detection, potentially dangerous situations may develop in pressure-containing applications. Environments Promoting Graphitic Corrosion (Ref 9). Experience has demonstrated that graphitic corrosion is favored by relatively mild environments such as soft waters, waters having a slightly acidic pH, waters containing low levels (as little as 1 ppm) of hydrogen sulfide, and brackish and other high-conductivity waters. It should not be inferred from this that gray or nodular cast irons are immune to more aggressive environments but, rather, that graphitic corrosion is less likely in them. In more aggressive environments, corrosion may indeed occur but may be manifested as general metal loss rather than as graphitic corrosion. Moist soils, especially those containing sulfates, will frequently produce graphitic corrosion of unprotected gray and nodular cast iron. Stray currents have also been identified as causes of graphitic corrosion in subterranean pipelines. Identification (Ref 10). Cast iron is converted to a soft mixture of iron oxides and graphite. Pieces of corrosion product smudge hands and
Forms of Corrosion: Recognition and Prevention
can be used to mark paper, just as if the corroded material were lead in a pencil. Attack is often uniform, with all exposed surfaces corroded to roughly the same depth (Fig. 51). If localized deposits are present, especially those containing sulfate and chloride or other acidic species, corrosion may be confined to pockets. When attack is severe or prolonged, the entire component is converted to corrosion product. Surface contour
(a)
(b)
Fig. 51
Examples of graphitic corrosion of gray cast-iron parts. (a) Cross sections cut through graphitically corroded regions will readily show bright, intact metal surrounded by a soft, dark, corroded area. (b) Graphitically corroded valve butterfly. Original surface contours are preserved. Edges, which were completely converted to brittle product, have broken. Source: Nalco Chemical Company
163
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Corrosion: Understanding the Basics
and appearance are often preserved. Attack is usually not apparent until surfaces are probed or stressed. Prevention. Attack by graphitic corrosion is reduced by the following: · Alloy substitution. Ductile cast iron is less prone to serious graphitic corrosion than gray irons, although it is not immune. White cast iron, which is essentially free of graphite, is immune to graphitic corrosion. Corrosion-resistant cast irons containing chromium, nickel, and silicon are also immune to graphitic corrosion. · Raising water pH to neutral or slightly alkaline levels to decrease attack, especially if relatively high concentrations of aggressive anions such as chloride and sulfate are present. · Use of inhibitors · Avoidance of stagnant water conditions
Stress-Corrosion Cracking General Description. Stress-corrosion cracking (SCC) is a cracking phenomenon that occurs in susceptible alloys and is caused by the conjoint action of a surface tensile stress and the presence of a specific corrosive environment. For SCC to occur on an engineering structure, three conditions must be met simultaneously, namely, a specific crack-promoting environment must be present, the metallurgy of the material must be susceptible to SCC, and the tensile stresses must be above some threshold value. Stresses required to cause SCC are small, usually below the macroscopic yield stress. The stresses can be exter-
Fig. 52
200 mm
Intergranular stress corrosion crack produced in 7050-T651 following exposure to 90 °C (195 °F), 90% relative humidity air. Specimens were etched in 10% NaOH at 70 °C (160 °F) for 20 s, nitric acid rinse.
Forms of Corrosion: Recognition and Prevention
165
nally applied, but residual stresses often cause SCC failures. This cracking phenomenon is of particular importance to users of potentially susceptible structural alloys because SCC occurs under service conditions that can result, often with no warning, in catastrophic failure. Stress-corrosion cracks often undergo extensive branching and proceed in a general direction perpendicular to the stresses contributing to their initiation and propagation. Figure 52 shows intergranular SCC in a specimen of high-strength aluminum alloy 7050 that occurred following exposure to 90 °C (195 °F), 90% relative humidity air. Figure 53 shows an SCC failure of a titanium alloy pressure vessel used to contain uninhibited nitrogen tetroxide (N2O4) under pressure. The penetration of multiple stress-corrosion cracks from the internal surface of the tank is visible. A group of these cracks eventually grew to such a depth that the remaining wall thickness was unable to support the pressure, and mechanical failure occurred. Figure 54 shows a type 304 stainless steel heat exchanger tube that failed after exposure to chlorides. The cracks were highly branched and transgranular. As these figures indicate, in some metals, cracking propagates intergranularly; in others, transgranularly. In certain metals, such as high-nickel alloy, iron-chromium alloys, and brasses, either type of cracking can occur, depending on the metal-environment combination.
(a) (b)
Fig. 53
Stress-corrosion failure of an Apollo Ti-6Al-4V RCS pressure vessel due to nitrogen tetroxide. (a) Failed vessel after exposure to pressurized N2O4 for 34 h. (b) Cross section through typical stress-corrosion cracks. 250´
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Corrosion: Understanding the Basics
Materials Factors Influencing SCC Behavior. The alloy composition and microstructure have a great effect on the susceptibility of a material to SCC in a particular environment. The bulk alloy composition may affect the formation and stability of a protective film on the surface. The alloy composition includes the nominal composition, the presence of constituents, and the presence and composition of impurities or trace elements. The metallurgical condition, which affects the susceptibility to SCC, includes the strength level, the presence of
(a)
(b)
Fig. 54
Stress-corrosion failure of a type 304 stainless steel heat exchanger tube from carbon dioxide compressor intercooler after exposure to a pressurized chloride-containing (200 ppm) environment at 120 °C (250 °F) (a) Cracks on the external surface. (b) Cracks originating on the external surface. Note the branching of the cracks. 33´. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
phases in the matrix and at the grain boundaries, the composition of the phases, the grain size and orientation of the phases, the grain size and orientation, grain-boundary segregation, and residual stresses. An example of strong influence of alloy composition and microstructure on the susceptibility to SCC is given by austenitic stainless steels, where chromium and molybdenum promote the formation of passive films on the surface. Trace elements such as carbon at concentrations greater than 0.03 wt%, may cause sensitization by forming chromium carbides at the grain boundaries and depleting zones around the carbides of chromium, thereby rendering the steel susceptible to intergranular SCC (IGSCC). Austenitic stainless steels will fail transgranularly in high-temperature chloride solutions. Similarly, the susceptibility of aluminum alloys to SCC strongly depends on the microstructure, which can be modified by heat treatment. The 7000 series aluminum alloys are precipitation-hardening alloys, and the peak-aged microstructure (T6) is the most susceptible to SCC. Overaging to the T76 or T73 condition usually reduces or eliminates the susceptibility to cracking. Peak aging of this alloy results in a fine distribution of coherent precipitates, which give strength to the alloy. However, the heat treatment also results in the formation of large incoherent precipitates at the grain boundaries and the depletion of solute in the region adjacent to the grain boundaries. Environmental Factors Influencing SCC Behavior. Stress-corrosion cracking of susceptible alloys is environment specific. NACE International and the Materials Technology Institute of the Chemical Process Industries and others have published tables of corrodents known to cause SCC of various metals and alloy systems. Table 7 lists these data in condensed form and covers the SCC environments of major importance to materials engineers. This table, as well as those published in the literature, should be used only as a guide for screening candidate materials for further in-depth investigation, testing, and evaluation. Stress-corrosion cracking is not a certainty in the listed environments under all conditions. Metals and alloys that are indicated as being susceptible can give good service under specific conditions (Table 7): · Anhydrous ammonia will cause SCC in carbon steels, but rarely at temperatures below 0 °C (32 °F) and only when such impurities as air or oxygen are present; addition of a minimum of 0.2% H2O will inhibit SCC. · Aqueous fluorides and hydrofluoric acid (HF) primarily affect Monel alloy 400 (UNS N04400) in the nickel alloys system; others are resistant. · Steam is known to cause SCC only in aluminum bronzes and silicon bronzes in the copper alloys system.
167
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Corrosion: Understanding the Basics
· Polythionic acid only cracks sensitized austenitic stainless steels and nickel alloys; SCC is avoided by solution annealing heat treatments or selection of stabilized or low-carbon alloys.
More detailed information on SCC of specific alloy systems can be found in Chapter 6. Environments that cause SCC are usually, but not necessarily, aqueous, and specific environmental parameters must be in specific ranges for cracking to occur. These include, but are not limited to, the following: · · · · · ·
Temperature pH Electrochemical potential Solute species Solute concentration Oxygen concentration
Changing any of these environmental parameters may significantly affect the crack nucleation process or the rate of crack propagation. Although the parameters listed above are important in controlling the rate of SCC, conditions inside a propagating crack and at the crack tip, which actually control the crack propagation process, are often quite Table 7
Some environment-alloy combinations known to result in SCC
Environment
Amines, aqueous Ammonia, anhydrous Ammonia, aqueous Bromine Carbonates, aqueous Carbon monoxide, carbon dioxide, water mixture Chlorides, aqueous Chlorides, concentrated, boiling Chlorides, dry, hot Chlorinated solvents Cyanides, aqueous, acidified Fluorides, aqueous Hydrochloric acid Hydrofluoric acid Hydroxides, aqueous Hydroxides, concentrated, hot Methanol plus halides Nitrates, aqueous Nitric acid, concentrated Nitric acid, fuming Nitrites, aqueous Nitrogen tetroxide Polythionic acids Steam Sulfides plus chlorides, aqueous Sulfurous acid Water, high-purity, hot X, known to result in SCC
Alloy system Stainless steels Austenitic Duplex
Aluminum alloys
Carbon steels
Copper alloys
Nickel alloys
¼ ¼ ¼ ¼ ¼ ¼
X X ¼ ¼ X X
X ¼ X ¼ ¼ …
¼ ¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ ¼ ¼ ¼
X ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ X
¼ ¼ ¼ ¼ X ¼ ¼ ¼ X ¼ ¼ X ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ ¼ ¼ ¼ ¼ … ¼ ¼ ¼ X ¼ ¼ X ¼ ¼ X ¼ ¼ ¼
X X X ¼ ¼ X ¼ X ¼ X ¼ ¼ ¼ ¼ ¼ ¼ X ¼ ¼ ¼ X
X X ¼ ¼ ¼ ¼ ¼ ¼ X X ¼ ¼ ¼ ¼ ¼ ¼ X ¼ X X ¼
Martensitic
Titanium alloys
Zirconium alloys
¼ ¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ ¼ ¼ ¼
¼ ¼ ¼ X ¼ ¼
X X ¼ ¼ ¼ ¼ ¼ ¼ X X ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ X ¼ ¼
¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼ X X ¼ X ¼ ¼ ¼ ¼ ¼ ¼ X ¼ ¼
¼ ¼ X X ¼ ¼ X ¼ ¼ ¼ X ¼ ¼ X ¼ X ¼ ¼ ¼ ¼ ¼
X ¼ ¼ X ¼ ¼ ¼ ¼ ¼ ¼ X ¼ X ¼ ¼ ¼ ¼ ¼ ¼ ¼ ¼
Forms of Corrosion: Recognition and Prevention
different from the so-called bulk environmental parameters. The pH inside cracks often differs from that in the bulk environment. In low-alloy steels containing approximately 1 wt% Cr, dissolution and hydrolysis of chromium can result in a lowering of the pH to values near 4 (Ref 11). In the case of stainless steels, the pH value in cracks can range from 0 to 3, with the lowest pH values associated with concentrated salt solutions containing chromium and ferrous ions. The pH inside cracks of aluminum and aluminum alloys is generally in the range of 3 to 4, while the pH inside propagating cracks in titanium alloys can be as low as 1 (Ref 11). Stress-corrosion cracking consists of a crack nucleation and propagation phase. Very little is known about the conditions that control the nucleation of a crack, other than that the thermodynamic and kinetic conditions must be right for the crack nucleate. For example, for anodically assisted SCC, metal dissolution and subsequent formation of a protective oxide film must be thermodynamically possible. The thermodynamic requirement of simultaneous dissolution and film formation has led to the identification of critical potentials at which SCC can occur. The thermodynamic conditions at which dissolution and film forming occurs is described by potential-pH (Pourbaix) diagrams. For example, the Pourbaix diagram in Fig. 55 describes the conditions at which dissolution and film formation on carbon steel can occur in different environments such as phosphate, nitrate, and carbonate/bicarbonate solutions. The effects of many of the external parameters listed above, such as pH, temperature, potential, and solute and oxygen concentration can have a great effect on thermodynamic stability and, thus, on the susceptibility to SCC. The diagram indicates that severe susceptibility to SCC is encountered when a protective film such as carbonate, phosphate, or magnetite is thermodynamically stable. In addition to the thermodynamic stability requirements for crack nucleation and propagation, kinetic requirements also need to be met. As in the thermodynamic requirements for SCC, environmental parameters such as potential, pH, solute and oxygen concentration, temperature, and crack-tip chemistry have a strong effect on the crack nucleation and crack growth kinetics. Figure 56 shows examples of potentio-dynamic polarization curves for alloy 600, alloy 800, and type 304 stainless steel in 10% NaOH solution at 288 °C (550 °F), indicating the various potential regions for susceptibility to SCC. The figure shows that, for all three alloys, the active-passive transition region presents the critical potential range for SCC to occur. However, the critical potentials, as well as the mode of cracking, are different for each alloy. Mechanical Factors Influencing SCC Behavior. Threshold stresses and stress-intensity factors, the presence of a stress-independent crack-growth regime, and the dependence of cracking to strain rate are important features in determining the susceptibility of alloys to SCC.
169
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Corrosion: Understanding the Basics
Potential, V vs. standard hydrogen electrode (SHE)
The threshold stress is typically the stress value obtained from constantload testing below which SCC does not occur and can serve as a simple measure for susceptibility of a material to SCC in a certain environment. The stress-intensity factor, K, is a parameter that describes the relationship between the applied stress and crack length for specific specimen geometries. Figure 57 shows the stress-intensity factor, K, as a function of the crack propagation rate, da/dt. The threshold is defined in this figure by the minimum detectable crack growth rate. The threshold stress intensity is generally associated with the development of a plastic zone at the crack tip. Stage I crack growth shows a rapid increase in crack growth rate, while in Stage II, the crack growth rate is independent of the stress intensity.
Fig. 55
Relationship between pH-potential conditions for SCC susceptibility of carbon steel in various environments and the stability regions for solid and dissolved species on the Pourbaix diagram. Note that severe susceptibility is encountered where a protective film (phosphate, carbonate, magnetite, and so on) is thermodynamically stable, but if ruptured, a soluble species (Fe 2 +, HFeO –2 ) is metastable.
Forms of Corrosion: Recognition and Prevention
The stress-corrosion crack propagation is usually studied with linear elastic fracture mechanics (LEFM), which assumes little plasticity at the tip of the propagating crack, such that the stress state is triaxial or plane strain. When the plastic zone size exceeds a certain value, either by increased stress, by propagation in a ductile material, or by crack propagation in a thin member, the stress state becomes biaxial or plane
Fig. 56
Log crack propagation rate, da/dt
Potentiokinetic polarization curve and electrode potential values at which intergranular and transgranular SCC appears in a 10% sodium hydroxide (NaOH) solution at 288 °C (550 °F). (a) Alloy 600. (b) Alloy 800. (c) AISI type 304 stainless steel
Fig. 57
Schematic of stress-corrosion crack velocity as a function of stress-intensity factor, K
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stress, and LEFM is not applicable. Then the more fundamental parameters, the energy release rate or J-integral, can be applied to describe the propagating stress-corrosion crack (Ref 12). The slow-strain-rate technique provides an excellent way to determine the susceptibility of an alloy to SCC (Ref 13, 14). However, the strain-rate behavior strongly depends on the alloy/environment combination. For example, for most materials, the critical strain rate, at which the maximum susceptibility is obtained, is 10–6/s. This critical strain rate points to a cracking mechanism whereby the rate of anodic dissolution is equal to the rate of protective film formation. If a higher strain rate is applied, the mechanical fracture will be more rapid than the rate of anodic dissolution. On the other hand, when a lower strain rate is applied, anodic dissolution will continue to blunt the crack, and SCC cannot occur. When other SCC mechanisms are predominant, the critical strain rate may be at a higher value, as is often the case with internal hydrogen embrittlement, or there may be no critical value, which occurs when the susceptibility decreases with decreasing strain rate. This has been observed in cases where the mechanism of SCC is thought to be hydrogen embrittlement. More detailed descriptions of the mechanisms of SCC and the use of LEFM for studying SCC behavior can be found in Corrosion, Volume 13, of the ASM Handbook.
Fig. 58
Grain orientations in standard wrought forms of alloys
Forms of Corrosion: Recognition and Prevention
Evaluation of SCC. In order to determine the susceptibility of alloys to SCC, several types of testing are available. If the objective of testing is to predict the service behavior or to screen alloys for service in a specific environment, it is often necessary to obtain SCC information in a relatively short period of time, which requires acceleration of testing by increasing the severity of the environment or the critical test parameters. The former can be accomplished by increasing the test temperature or the concentration of corrosive species in the test solution and by electrochemical stimulation. Test parameters that can be changed to reduce the testing time include the application of higher stresses, continuous straining, and precracking, which allows bypassing of the crack nucleation phase of the SCC process. Stress-corrosion specimens can be divided into two main categories, namely smooth, and precracked or notched specimens. Further distinction can be made in the loading mode, such as constant deflection, constant load, and constant extension or strain rate. During alloy processing, operations used in the production of wrought alloys, the metal is forced in a predominant direction so that the grains are elongated in the direction of flow. Because it is important to relate the application of stress and the grain flow direction, two conventions are used to relate the two parameters. In one system, which is primarily used for smooth specimens, the three stressing directions are designated by indicating the direction of the stress, namely longitudinal (L), long-transverse (LT), transverse (T), and short-transverse (Fig. 58). A second system, which is particularly useful for precracked specimens, indicates both the cracking plane and the direction of crack propagation. The system uses three letters (L, T, and W) to indicate three perpendicular directions, namely L for the longitudinal direction, T for the thickness direction, and W for the width direction. The crack plane is indicated by the direction normal to the crack, and the crack propagation is indicated by one of the directions L, T, or W. Figure 59 demonstrates the various orientations for a double-cantilever-beam (DCB)
Fig. 59
Fracture plane identification in double-cantilever-beam specimens. L, direction of grain flow; T, transverse grain direction; S, short transverse grain direction; C, chord of cylindrical cross section; R, radius of cylindrical cross section; first letter, normal to the fracture plane; second letter, direction of crack propagation in fracture plane
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specimen. More detailed information on SCC test specimens and test methods can be found in Ref 15. Prevention of SCC can be accomplished by the following: · Changing the material: A common method to control SCC is to select a material with greater SCC resistance to a particular environment. In other cases, the cracking may result from a specific microstructural condition of the metal (e.g., sensitized austenitic stainless steel). In this case, avoiding the susceptible microstructure can solve the problem. Often, as the strength of an alloy increases, its susceptibility to SCC increases. Use of alloys with lower strength levels can thus be an effective means of reducing the likelihood of SCC if the design/ application permits. · Changing the environment: The environment can be modified to reduce the concentration of aggressive species or to reduce the temperature. Inhibitors can also be added to control SCC. · Changing the oxidizing potential: Shifting the electrochemical potential of the metal’s surface out of the critical potential range for cracking will prevent SCC. The potential can be shifted to more oxidizing potentials (anodic protection) or to more reducing potentials (cathodic protection).
Fig. 60
Methods used to control SCC. Source: Ref 16
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· Applying protective coatings: Coatings that isolate the metal substrate from the corrosive environment can effectively control SCC. In these cases, however, the protection depends on the integrity of the protective film. · Changing applied or residual stress levels: Several approaches can be used to alleviate SCC problems. The level of tensile stresses at the metal surface can be lowered either by lowering the operating stresses, stress relieving the parts subsequent to fabrication to reduce residual stresses, or shot peening to build up compressive stresses at the metal surface. Another important consideration is to use proper grain orientations to reduce sustained stress levels on the part below the stresscorrosion threshold levels, especially in the more susceptible shorttransverse direction grain directions. High-strength aluminum alloys have much lower SCC resistance in the short-transverse direction.
Figure 60 summarizes the various approaches to controlling SCC.
Corrosion Fatigue General Description. Corrosion fatigue is a term that is used to describe the phenomenon of cracking, including both initiation and propagation, in materials under the combined actions of a fluctuating, or cyclic stress, and a corrosive environment. Corrosion fatigue depends strongly on the interactions between the mechanical (loading), Table 8 Mechanical, metallurgical, and environmental variables that influence corrosion fatigue behavior Variable
Type
Mechanical
Maximum stress or stress-intensity factor, smax or Kmax Cyclic stress or stress-intensity range, Ds or DK Stress ratio (R) Cyclic loading frequency Cyclic load waveform (constant-amplitude loading) Load interactions in variable-amplitude loading State of stress Residual stress Crack size and shape, and their relation to component size and geometry
Metallurgical
Alloy composition Distribution of alloying elements and impurities Microstructure and crystal structure Heat treatment Mechanical working Preferred orientation of grains and grain boundaries (texture) Mechanical properties (strength, fracture toughness, etc.)
Environmental
Temperature Types of environments: gaseous, liquid, liquid metal, etc. Partial pressure of damaging species in gaseous environments Concentration of damaging species in aqueous or other liquid environments Electrical potential pH Viscosity of the environment Coatings, inhibitors, etc.
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metallurgical, and environmental variables listed in Table 8. As shown in Fig. 61, an aggressive environment usually has a deleterious effect on fatigue life, producing failure in fewer stress cycles than would be required in a more inert environment. Corrosion fatigue produces fine-to-broad cracks with little or no branching (Fig. 62 and 63); thus, they differ from stress-corrosion cracks, which often exhibit considerable branching. They are typically filled with dense corrosion product (Fig. 63a). The cracks may occur singly but commonly appear as families or parallel cracks. They are frequently associated with pits, grooves, or some other form of stress concentrator. Transgranular fracture paths are more common than intergranular fractures. Prevention. All metals and alloys are susceptible to corrosion fatigue. Even some alloys that are immune to SCC, for example, ferritic stainless steels, are subject to failure by corrosion fatigue. Both temporary and permanent solutions for corrosion involve reducing or eliminating cyclic stresses, selecting a material or heat treatment with higher corrosion fatigue strengths, reducing or eliminating corrosion, or a combination of these procedures. These objectives are accomplished by changes in material, design, or environment.
Fig. 61
Typical fatigue behavior in an aggressive environment compared with fatigue behavior in an inert environment or at high frequency. (a) Data plotted as S-N curves. (b) Data plotted as the crack-growth rate vs. stress-intensity range. The environmental influence is most pronounced at low and intermediate stress-intensity levels. At high stress intensity, the mechanical crack-growth driving force becomes dominant, and the influence of environment is less significant. As the stress intensity approaches the fracture toughness of the material where instability occurs, the effect of the aggressive environment becomes less significant
Forms of Corrosion: Recognition and Prevention
For permanent solutions, if a new material is selected (the same alloy with a different heat treatment or fabrication method, another alloy for the same system, or an alloy from a completely different system), basic compatibility with design requirements and with the environment to be encountered in service usually must be proved by testing. If changes in design or in mode of fabrication are selected, test specimens should reflect these changes. Any environmentally affected properties must be checked in a simulated service environment. Any alteration of the environment should also be testing at this stage. Finally, the ultimate test of a redesigned component is its service life in comparison with the life of the component that is replaced. A review of corrosion fatigue test methods can be found in the article “Corrosion Fatigue Testing” in Volume 19 of the ASM Handbook. Operating stress may be lowered by reducing either the mean stress or the amplitude of the cyclic stress. This almost always involves a change in component design. Sometimes, only a minor change is required, such as increasing a fillet radius to reduce the amount of stress concentration at a critical location. In other instances, more extensive changes are required, such as significantly increasing the cross-sectional area or adding a strengthening rib. If failure has occurred because of stress raisers introduced during manufacture, changes in manufacturing processes or procedures and quality requirements may be necessary.
Fig. 62
Corrosion-fatigue cracks in carbon steel. A nital-etched section through corrosion-fatigue cracks that originated at hemispherical corrosion pits in a carbon steel boiler tube. Corrosion products are present along the entire length of the cracks. 250×
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(a)
(b)
Fig. 63
Failure of boiler tube wall due to corrosion fatigue cracking. (a) Wedge-shaped corrosion fatigue crack filled with corrosion product. As the cyclic process continues, this crack will eventually propagate through the tube wall. (b) A family of longitudinal corrosion fatigue cracks resulting from fluctuations in internal pressure. Source: Nalco Chemical Company
Forms of Corrosion: Recognition and Prevention
Shot peening is usually effective in prolonging fatigue life in air by introducing residual compressive stresses in a metal surface, thus reducing the mean tensile stress at potential crack-initiation sites. In more aggressive environments, however, shot peening may have only limited value, because general corrosion can eventually remove the surface layer and, thus, the beneficial compressive residual stresses. Nitriding, which also introduces compressive residual surface stresses, can improve the corrosion-fatigue resistance of steels, particularly when a relatively short life is required. For applications requiring extended life, other measures are more effective. Material strength is generally increased by alloying, heat treatment, cold working, or selection of a material from a different alloy system. Because corrosion may reduce the fatigue strength of a material to only a small fraction of its strength in air, alloying additions or heat treatments that only increase strength, without altering the corrosion resistance of the material, may be of marginal value. Conversely, alloying that improves corrosion resistance can be effective in combating corrosion fatigue. For instance, chromium-containing steels are generally more corrosion-fatigue resistant than carbon steels. Galvanic protection by sacrificial anodes or applied cathodic currents has been successful in reducing the influence of corrosion on fatigue of metals and alloys exposed to aqueous environments, except in such alloys as high-strength steels that are subject to hydrogen-induced delayed cracking. In similar instances, anodic polarization has been used for protection of stainless steels. With passive alloys or alloys that can be polarized to produce passive behavior, crevices must be avoided because corrosion within crevices may actually proceed more rapidly due to the anode-cathode relationship. For example, in a part made of low-carbon steel, an area under an O-ring can fail by corrosion fatigue in a caustic solution with a pH of 12 at a location where only low cyclic stresses exist. Where pitting corrosion has occurred, corrosion pits may act as stress raisers and thus accelerate fatigue failure. Inhibitors are sometimes added to the environment or included in organic coatings to eliminate corrosion fatigue. The effect of inhibitors is believed to depend solely on their ability to reduce corrosion rates to acceptable values. Corrosion fatigue cannot occur if there is no contact between the surface of a susceptible material and a corrosive environment; however, ordinary fatigue can occur. In most instances, the environment cannot be prevented from contacting the component; therefore, coatings are necessary. Continuity of the coating is important. Organic coatings, such as paint or plastic, are only physical barriers (unless they contain inhibitors) and therefore must be absolutely continuous to be effective. The density and the thickness of coatings are also significant factors in the prevention of corrosion fatigue. Noble-metal coatings can be effec-
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tive, but only if they remain unbroken and are of sufficient density and thickness. The relatively low corrosion-fatigue strength of carbon steel is reduced still further when local breaks in a coating such as nickel occur. Presumably, noble-metal coatings that contain residual compressive stresses are more effective than coatings that contain residual tensile stresses.
Hydrogen Damage The term hydrogen damage has been used to designate a number of processes in metals by which the load-carrying capacity of the metal is reduced due to the presence of hydrogen, often if combination with residual or applied tensile stresses. Although it occurs most frequently in carbon and low-alloy steels, many metals and alloys are susceptible to hydrogen damage. Hydrogen damage in one form or another can severely restrict the use of certain materials. Because hydrogen is one of the most abundant elements and is readily available during the production, processing, and service of metals, hydrogen damage can develop in a wide variety of environments and circumstances. The interaction between hydrogen and metals can result in the formation of solid solutions of hydrogen in metals, molecular hydrogen, gaseous products that are formed by reactions between hydrogen and elements constituting the alloy, and hydrides. Depending on the type of hydrogen/metal interaction, hydrogen damage of metal manifests itself in one of several ways. Specific types of hydrogen damage, some of which occur only in specific alloys under specific conditions are the following: · · · · ·
Hydrogen embrittlement Hydrogen-induced blistering Cracking from precipitation of internal hydrogen Hydrogen attack Hydride formation
The first three types are usually observed at ambient temperatures and are closely related to one another. Hydrogen attack is an elevatedtemperature phenomenon that affects carbon and alloy steels. Hydride formation produces embrittlement and cracking in transition metals such as titanium, tantalum, and zirconium.
Hydrogen Embrittlement General Description. When high-strength steel containing hydrogen is stressed in tension, even if the applied stress is less than the yield
Forms of Corrosion: Recognition and Prevention
181
strength, it may fail prematurely in a brittle manner. This type of hydrogen damage occurs most often in high-strength steels, primarily quenched-and-tempered steels and precipitation-hardened steels. The presence of hydrogen in steel reduces the tensile ductility and causes premature failure under static load that depends on the stress and time. This phenomenon is known as hydrogen embrittlement. Steel can be embrittled by a very small amount of hydrogen, often a few parts per million, and hydrogen may come from various sources (e.g., acid pickling, electroplating, and aqueous corrosion), all of which are electrochemical processes involving the discharge of hydrogen ions. Unlike SCC, cracks caused by hydrogen embrittlement usually do not branch. Failure by the hydrogen embrittlement mechanism is accompanied by very little plastic deformation, and the fracture mode is often brittle intergranular fracture (Fig. 64).
(a)
(c)
(b)
Fig. 64
Hydrogen embrittlement failure of a 300 M steel space shuttle orbiter nose landing gear steering collar pin. The pin was heat treated to a 1895-MPa (275 ksi) strength level. The part was plated with chromium and titanium-cadmium. (a) Pin showing location of failure (actual size). (b) Failure origin (arrow). 9×. (c) Brittle intergranular fracture face characteristic of hydrogen embrittlement. Parts did not receive a hydrogen embrittlement relief bake due to processing error. 1380×
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Cracking from Hydrogen Charging in an Aqueous Environment. When metal corrodes in a low-pH solution, the cathodic partial reaction is reduction of the hydrogen ion. Although most of the reduced hydrogen reacts to form H2 and leaves the metal surface as gaseous hydrogen, a part of the reduced hydrogen enters into the metal as atomic hydrogen. The presence of certain chemical substances that prevent the recombination of hydrogen to form molecular hydrogen enhances the absorption of nascent (atomic) hydrogen into the metal. These substances are called cathodic poisons, and they include phosphorus, arsenic, antimony, sulfur, selenium, tellurium, and cyanide ion. Among the cathodic poisons, sulfide is the most common. Environments containing hydrogen sulfide can cause severe embrittlement of steels and some other high-strength alloys. On the other hand, corrosion inhibitors lower the corrosion rate and thus the amount of hydrogen charged into the metal. Atmospheric and Aqueous Corrosion. Most high-strength steels are susceptible to hydrogen embrittlement when stressed and exposed to fresh or seawater and even during atmospheric exposure. The susceptibility of steels to hydrogen embrittlement generally increases with increasing tensile strength. Steels having a tensile strength greater than about 1034 MPa (150 ksi) are much more susceptible to embrittlement. Above a tensile-strength level of 1241 MPa (180 ksi), most high-strength low-alloy steels, such as AISI 4130 and 4340, and precipitation-hardening stainless steels are susceptible to hydrogen-embrittlement cracking in marine atmospheres when the residual or applied tensile stresses are sufficiently high, and the cracking usually occurs in a form of delayed failure. Steels with tensile strengths less than 690 MPa (100 ksi) appear tpo be resistant to hydrogen-embrittlement cracking, and the structures made with such steels have been used in service without serious problems in various environments that do not contain hydrogen sulfide. Environments Containing Hydrogen Sulfide. High-strength steel pipes used in drilling and completion of oil and gas wells may exhibit delayed failure in environments containing hydrogen sulfide. This type of failure is referred to as sulfide stress cracking. The basic cause of sulfide stress cracking is embrittlement resulting from hydrogen absorbed into steel during corrosion in sour environments. The presence of hydrogen sulfide in the environment promotes hydrogen absorption into steel, thereby making the environment more severe and thus more likely to cause hydrogen embrittlement. Although hydrogen sulfide gas, like gaseous hydrogen, can cause embrittlement, water ordinarily must be present for sulfide stress cracking to occur. The susceptibility to sulfide stress cracking increases with increasing hydrogen sulfide concentration or partial pressure and decreases with increasing pH. The ability of the environment to cause sulfide stress cracking decreases markedly above pH 8 and below 101 Pa (0.001 atm)
Forms of Corrosion: Recognition and Prevention
partial pressure of hydrogen sulfide. The cracking tendency is most pronounced at ambient temperature and decreases with increasing temperature. For a given strength level (commonly measured nondestructively by hardness), tempered martensitic steels have better sulfide stress-cracking resistance than normalized-and-tempered steels, which in turn are more resistant than normalized steels. Untempered martensite demonstrates poor resistance to sulfide stress cracking. It is generally agreed that a uniform microstructure of fully tempered martensite is desirable for sulfide stress-cracking resistance. In most cases, carbon and low-alloy steels are used at hardnesses of 22 HRC or below to prevent sulfide stress cracking. The effect of alloying elements on the sulfide stress-cracking resistance of carbon and low-alloy steels is controversial, except for one element. Nickel is detrimental to sulfide stress-cracking resistance. Steels containing more than 1% Ni are not recommended for service in sour environments. Pickling and Electroplating. Pickling alone is not a serious problem (unless internal voids or other imperfections lead to local formation of molecular hydrogen) because the metal is not being stressed, and much of the absorbed hydrogen diffuses out of the metal. Heating at 150 to 200 °C (300 to 390 °F) for 3 h hastens the removal of hydrogen. Also, the addition of suitable inhibitors to the pickling solution eliminates or minimizes attack on the metal and the consequent generation of nascent hydrogen. Salt baths operated at about 210 °C (410 °F) can be used for descaling titanium alloys, superalloys, and refractory metals to avoid the possibility of hydrogen charging associated with pickling. Plating solutions and plating conditions selected to produce a high-cathode efficiency minimize the amount of hydrogen generated on the metal. Because the metallic coatings plated on metal often act as barriers to effusion of hydrogen, elevated-temperature baking after plating is generally required for removal of hydrogen. Baking at 190 °C (375 °F) usually suffices, unless the coating is cadmium, through which hydrogen diffuses less readily than through other electrodeposited metals. Raising the baking temperature to accelerate effusion is not possible without reducing the protective qualities of the cadmium plate. Thus, baking time must be lengthened, particularly for steels with tensile strengths exceeding 1380 MPa (200 ksi). Baking times up to 24 h are necessary for such steels. Cracking from Gaseous Hydrogen. Steel vessels and other equipment that contains hydrogen gas at high pressures at ambient temperature are susceptible to failure by hydrogen embrittlement. Many high-strength steels are severely embrittled in tension tests conducted in high-pressure hydrogen gas, and the cracking tendency increases with increasing pressure of hydrogen. The cracking susceptibility of steels generally increases with increasing strength.
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Generally, high-strength steels and high-strength nickel alloys display severe degradation of tensile properties in hydrogen gas. Austenitic stainless steels, aluminum alloys, and alloy A-286 show very little embrittlement in this environment. Most other engineering metals and alloys are affected to a lesser degree.
Hydrogen-Induced Blistering General Description. Hydrogen-induced blistering, also referred to as hydrogen-induced cracking (HIC) or blister cracking, is primarily found in lower-strength (unhardened) steels, typically with tensile strengths less than approximately 550 MPa (80 ksi). It is observed in steels that have been exposed to hydrogen-charging conditions, for example, pickling or corrosion in environments containing hydrogen sulfide. During such exposures, atomic hydrogen generated at the metal surface is absorbed by the metal. When hydrogen is absorbed into the metal and diffuses inward, it can precipitate as molecular hydrogen at internal voids, laminations, or inclusion/matrix interfaces, and it can build up pressure great enough to produce internal cracks. If these cracks are just below the surface, the hydrogen-gas pressure in the cracks can lift up and bulge out the exterior layer of the metal so that it resembles a blister (Fig. 65). The equilibrium pressure of the molecular hydrogen in the void, which is in contact with the atomic hydrogen in the surrounding metal, its great enough to rupture any metal or alloy. Hydrogen-Induced Cracking of Line-Pipe Steel. Line pipe transmitting wet, sour gas can develop hydrogen-induced cracking in the pipe wall. When line-pipe steel is exposed to sour brine in the absence of applied stress, a number of cracks parallel to the longitudinal axis of
Fig. 65
Hydrogen-induced blistering in a 9.5 mm (3 8 in.) thick carbon steel (ASTM A 285 Grade C) plate that had been in service 1 year in a refinery vessel. 1 12×
Forms of Corrosion: Recognition and Prevention
the pipe may develop through the wall, and the tip of one crack may link up with another in stepwise fashion. This type of cracking is called stepwise cracking (Fig. 66), and it can significantly reduce the effective wall thickness of the pipe. It is believed that a single, straight longitudinal crack is less harmful than stepwise cracks. The presence of hydrogen sulfide in the corrodent greatly promotes the absorption of hydrogen into steel. Factors That Affect HIC. Aside from reducing the amount of hydrogen being generated by reducing the corrosion reaction, another way of controlling HIC is through material processing. Shape control of sulfide inclusions is perhaps the best way to minimize the tendency toward HIC in line-pipe steels. Elongated manganese sulfide inclusions promote crack initiation and propagation due to the high stresses at the tips of the inclusions. However, the addition of calcium or rare earths to the steel makes the sulfides spherical, and because of their hardness, they remain spherical after processing. In addition, reduction of the sulfur content is also beneficial in reducing the susceptibility of steels to HIC. Other alloying additions that reduce hydrogen permeation, such as copper up to about 0.25%, are also beneficial.
Cracking from Precipitation of Internal Hydrogen General Description. Shatter cracks, flakes, and fish eyes are features common to hydrogen damage in forgings, weldments, and castings. They are attributed to hydrogen pickup during melting operations when the melt has a higher solubility for hydrogen than the solid alloy. During cooling from the melt, hydrogen diffuses to and precipitates in voids and discontinuities, producing the features that result from the decreased solubility of hydrogen in the solid metal. In many aspects, these features are comparable to blistering, and this could be considered a special case of that class.
Fig. 66
Stepwise cracking of a low-strength pipeline steel exposed to hydrogen sulfide (H2S). 6×
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Hydrogen Attack General Description. Steel exposed to high-temperature, high-pressure hydrogen appears to be unaffected for days or months and then suddenly loses its strength and ductility. This type of damage is called hydrogen attack. It is important to note that hydrogen attack is different from hydrogen embrittlement. Hydrogen attack is irreversible damage, and it occurs at elevated temperatures, whereas hydrogen embrittlement is often reversible and occurs at temperatures below 200 °C (390 °F). In hydrogen attack of steel, absorbed hydrogen reacts internally with carbides to produce methane bubbles along grain boundaries. These bubbles subsequently grow and merge to form fissures. Failure by hydrogen attack is characterized by decarburization and fissuring at grain boundaries (Fig. 67) or by bubbles in the metal matrix. This type of hydrogen damage is most commonly experienced in steels that are subjected to elevated temperatures in petrochemicals-plant equipment that often handles hydrogen and hydrogen-hydrocarbon streams at pressures as high as 21 MPa (3 ksi) and temperatures up to 540 °C (1000 °F). The severity of hydrogen attack depends on temperature, hydrogen partial pressure, stress level, exposure time, and steel composition. Prevention of Hydrogen Attack. The only practical way to prevent hydrogen attack is to use only steels that, based on plant experience, have been found to be resistant to this type of deterioration. The following general rules are applicable to hydrogen attack: · Carbide-forming alloying elements, such as chromium and molybdenum, increase the resistance of steel to hydrogen attack.
(a)
(b)
Fig. 67
Section of ASTM A 106 carbon steel pipe with wall severly damaged by hydrogen attack. The pipe failed after 15 months of service in hydrogen-rich gas at 34.5 MPa (5000 psig) and 320 °C (610 °F). (a) Overall view of failed pipe section. (b) Microstructure of hydrogen-attacked pipe near the midwall. Hydrogen attack produced grain-boundary fissures that are radially aligned.
Forms of Corrosion: Recognition and Prevention
· Increased carbon content decreases the resistance of steel to hydrogen attack. · Heat-affected zones are more susceptible to hydrogen attack than the base or weld metal.
For most refinery and petrochemical plant applications, low-alloy chromiumand molybdenum-containing steels are used to prevent hydrogen attack. However, questions have recently been raised regarding the effect of long-term hydrogen exposure on C-0.5Mo steel. As a result, low-alloy steels are preferred over C-0.5Mo steel for new construction. The conditions under which different steels can be used in hightemperature hydrogen service are listed in the American Petroleum Institute (API) document 941 “Steels for Hydrogen Service at Elevated Temperatures and Pressures in Petroleum Refineries and Petrochemical Plants.” The principal data are presented in the form of Nelson curves, as shown in Fig. 68. The curves are periodically revised by the API Subcommittee on Materials Engineering and Inspection, and the latest edition of API 941 should be consulted to ensure that the proper steel is selected for the operating conditions encountered.
Hydride Formation General Description. A number of transition, rare-earth alkalineearth metals, and the alloys of these metals are subject to embrittlement and cracking due to hydride formation. Among these metals, the commercially important ones are titanium, tantalum, zirconium, uranium,
Fig. 68
Operating limits for various steels in high-temperature, high-pressure hydrogen service (Nelson curves) to avoid decarburization and fissuring
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Corrosion: Understanding the Basics
Fig. 69
Micrograph of severely hydrided unalloyed titanium. Approximately 200×
thorium, and their alloys. The presence of hydrides in these metals can cause significant increases in strength and large losses in ductility and toughness. As in other types of alloys, excess hydrogen is readily picked up during melting or welding, and hydride formation takes place during subsequent cooling. The use of vacuum melting and the modification of compositions can reduce susceptibility to hydride formation. Hydrogen can often be removed by annealing in vacuum. Welding generally requires the use of inert-gas shielding to minimize hydrogen pickup. The hydride particles often have the form of platelets and show preferred orientation within the parent lattice (Fig. 69), depending primarily on the metal or alloy composition. The large volume change associated with hydride formation leads to a strong interaction between the hydride-formation process and externally applied stresses. Applied stresses can cause preferential alignment of hydrides or realignment. In most cases, the hydride phase has a much lower ductility than the matrix.
Prevention of Hydrogen Damage The primary factors controlling hydrogen damage are material, stress, and environment. Hydrogen damage can often be prevented by using more resistant material, changing the manufacturing processes, modifying the design to lower stresses, or changing the environment. As described earlier, inhibitors and post-processing bake-out treatments can also be used. Selection of materials resistant to hydrogen damage is often possible. In many applications, a lower-strength material will function just as
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189
well mechanically as will a higher-strength material, and the use of such a material may eliminate a hydrogen-embrittlement problem. Coatings and linings can be applied to the surface of the metal to shield it from the environment. Such coating materials should be stable in the environment. Tempering. In high-strength steels, tempering at a higher temperature for a longer time lowers the strength and improves the toughness, and such a heat treatment would also help to reduce hydrogen content. Although such tempering treatments can improve resistance to hydrogen embrittlement, care should be exercised to avoid temper embrittlement. Temper-embrittled materials are more prone to cracking. Baking of electroplated parts reduces the possibility of hydrogen embrittlement. Surface preparation techniques that impart residual compressive stresses to the surface are used to improve resistance to cracking when susceptible metals are subjected to the combination of hydrogen embrittlement and residual or applied tensile stress. These techniques include shot peening, grit blasting, and face milling. Low-hydrogen welding rods should be specified for welding if hydrogen embrittlement is a potential problem. Also, it is important to maintain dry conditions in storing the rods before use and during welding because water and water vapor are major sources of hydrogen. Preheating and postweld heat treatment are also important when welding high-strength steels to prevent high residual stresses or microcracks. Lowering the stress below the threshold value for hydrogen embrittlement is often possible. This can be achieved by increasing the cross section of parts, avoiding stress raisers in design, reducing residual stresses through heat treatment, or reducing the load. Corrosion inhibitors reduce corrosion and therefore the amount of hydrogen absorbed into the metal, and they often eliminate the cracking problem. Proper use of inhibitors in the service environment can prevent hydrogenembrittlement failure.
Liquid-Metal Embrittlement Liquid-metal embrittlement (LME), also known as liquid-metal cracking, is the catastrophic brittle failure of a normally ductile metal when coated with a thin film of a liquid metal and subsequently stressed in tension. Although LME is not a corrosion process or even a metal dissolution process (like that encountered when a solid metal is immersed at length in a molten metal), it is included here because it is a problem frequently encountered by materials engineers. During
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LME, metals and alloys are penetrated, usually along grain boundaries, by such metals as mercury, which are liquid at room temperature, or by metals that have relatively low melting points. Examples of low-melting point metals contributing to LME are bismuth, tin, lead, cadmium, zinc, aluminum, and copper. Stress, temperature, and time are the factors that facilitate and accelerate LME. Virtually all metal and alloy systems are subject to LME by one or more of these metals at or above their melting points. More detailed information on LME can be found in Failure Analysis and Prevention, Volume 11 of the ASM Handbook (refer to p 225 to 244). Zinc is a prime offender because of widespread use throughout industry in the form of corrosion-resistant coatings applied to carbon steels by hot-dip galvanizing, electroplating, tumbling, and spray painting. Plain carbon steels are embrittled by zinc at temperatures above 370 °C (700 °F) for long periods of time, especially when the steel is heavily stressed or cold worked. Austenitic stainless steels and nickel-based alloys will also crack in the presence of molten zinc. These alloys usually crack instantly when welded to galvanized steel, a fairly common occurrence in the chemicalprocessing industry. In addition, austenitic alloy failures have occurred under the following conditions: · In high-temperature bolting fastened with galvanized steel nuts · During welding or heat treating of components contaminated by grinding with zinc-loaded grinding wheels, contact with zinc-coated structurals or slings, or exposure to zinc paint overspray · During process industry plant fires involving piping and vessels (thin-wall expansion joint bellows are especially susceptible) sprayed with molten zinc from coated steel structures
Thus, it is imperative that all traces of zinc be removed from coated steel members before welding to austenitic alloys and before intimate contact with these alloys at temperatures above 370 °C (700 °F). Also, austenitic stainless steels and nickel-based alloys should be handled with non-coated steel hoist chains, cables, and structurals. They should be dressed and cleaned with new grinding wheels and stainless steel brushes, and they should be marked with materials (paints, crayons, and so on) free from zinc and other low-melting metals. Cadmium is probably second to zinc in importance as an agent of liquid-metal embrittlement because of its applications as a corrosionresistant coating to a variety of hardware, particularly fasteners. Failures by cadmium LME of bolting operating at temperatures above 300 °C (570 °F) and fabricated from high-strength alloy steels, such as AISI 4140 and 4340, and austenitic stainless steels are fairly common. In fact, some high-strength steels and high-
Forms of Corrosion: Recognition and Prevention
191
strength titanium alloys are embrittled by cadmium at temperatures below its melting point by mechanisms not yet understood. The solution to LME by cadmium is similar to that of zinc, that is, avoidance of contact with, and contamination of, susceptible metal and alloy systems at temperatures above the 321 °C (610 °F) melting point of cadmium (and at all temperatures at which high-strength steels and titanium alloys are involved). Metal systems that are embrittled by contact with mercury include copper and its alloys, aluminum and its alloys, Nickel 200 (at elevated temperatures) and Monel Alloy 400, and titanium and zirconium and their alloys. Cracking is intergranular except in zirconium alloys; in these alloys, cracking is transgranular. Mercury LME of aluminum and copper alloys was more common years ago in the petrochemical industry when mercury-filled manometers and thermometers were extensively used. Failures or upsets would release mercury into process or service (steam, cooling water, and so on) streams, causing widespread cracking of piping, heat exchanger tube bundles, and other equipment. Under these conditions, even pure aluminum and pure copper are susceptible. With regard to the titanium system, the commercially pure grades used in the chemical-processing industry are less sensitive than the alloys. In addition, LME in aqueous solutions of mercurous salts, such as mercurous nitrate, is possible because the mercurous ion can be reduced to its elemental form at local cathodic sites.
References
1. M.G. Fontana, Eight Forms of Corrosion, Corrosion Engineering, 3rd ed., McGraw-Hill, Inc., 1986, p 67 2. A.J. Sedriks, Crevice Corrosion, Corrosion of Stainless Steels, 2nd ed., John Wiley & Sons, 1996, p 178 3. H.M. Herro and R.D. Port, Crevice Corrosion, The Nalco Guide to Cooling Water System Failure Analysis, McGraw-Hill, Inc., 1993, p 39–40 4. H.M. Herro, Paper 84, presented at Corrosion ’91, NACE International 5. L.S. Van Delinder, Ed., Localized Corrosion, Corrosion Basics: An Introduction, NACE International, 1984, p 97–98 6. R. Baboian, Phorgotten Phenomena: Galvanic Series Can Mislead, Mater. Perform., Aug 1998, p 70–71 7. R.A. Corbett, Phorgotten Phenomena: Galvanic Corrosion Can Occur At Same-Metal Couple, Mater. Perform., Dec. 1998, p 63–64
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8. H.M. Herro and R.D. Port, Cavitation Damage, The Nalco Guide to Cooling Water System Failure Analysis, McGraw-Hill, Inc. 1993, p 270–271 9. H.M. Herro and R.D. Port, Graphitic Corrosion, The Nalco Guide to Cooling Water System Failure Analysis, McGraw-Hill, Inc., 1993, p 376 10. R.D. Port and H.M. Herro, Graphitic Corrosion, The Nalco Guide to Boiler Failure Analysis, McGraw-Hill, Inc., 1991, p 261–263 11. A. Turnbull, Chemistry within Localized Corrosion Cavities, Advances in Localized Corrosion, June 1987 (Orlando, FL), NACE-9, NACE International, 1990 12. J.A. Begley and J.D. Landes, Proc. 1971 National Symposium on Fracture Mechanisms, Part III, STP, 514, ASTM, 1972, p 1 13. G.M. Ugianski and J.H. Payer, Ed., Stress-Corrosion Cracking— The Slow Strain-Rate Technique, STP 665, ASTM, 1979 14. J.A. Beavers and G.H. Koch, “Limitations of the Slow Strain Rate Test for Stress-Corrosion Cracking,” Publication 39, Materials Technology Institute of the Chemical Process Institute (MTI), 1995 15. D.O. Sprowls, Evaluation of Stress-Corrosion Cracking, Corrosion, Vol 13, ASM Handbook, ASM International, 1987, p 245–282 16. R.N. Parkins, An Overview—Prevention and Control of StressCorrosion Cracking, Mater. Perform., Vol 24, 1995, p 9–20
Selected References · Corrosion, Vol 13, ASM Handbook, ASM International, 1987, p 77–189 · C.P. Dillon, Ed., Forms of Corrosion—Recognition and Prevention, NACE Handbook 1, Vol 1, NACE International, 1982 · D. McIntyre, Ed., Forms of Corrosion—Recognition and Prevention, NACE Handbook 1, Vol 2, NACE International, 1997 · M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill, 1986
Corrosion: Understanding the Basics J.R. Davis, editor, p193-236 DOI: 10.1361/cutb2000p193
CHAPTER
Copyright © 2000 ASM International® All rights reserved. www.asminternational.org
5
Types of Corrosive Environments CORROSIVE ENVIRONMENTS can be broadly classified as atmospheric, underground/soil, water, acidic, alkaline, and combinations of these. Complicating matters is the fact that there are important variables, for example, pH, temperature, and the presence of biological organisms, that can significantly alter the response of the material in a given environment. Take, for example, the complex interaction of various elements that make up the chemistry of the environments in which automobiles are driven. The combination of natural environments (rain, snow, humidity, marine environments, etc.) with deleterious man-made environmental contributions (road salts, atmospheric pollutants and emissions, and acid rain) provide a uniquely aggressive corrosive environment for automobiles. The goal of this chapter is to point out the connection between the characteristics of the corrosive environment, the corrosion characteristics of various metals and materials systems, and the subsequent corrosion response. It should be noted that the corrosion literature contains a wealth of laboratory and service data on material/environmental combinations. In addition, ASM International and other associations, such as NACE International, publish useful information on corrosion in environments relevant to buildings, highways, and other structures, and components and equipment used in the chemicals processing, pulp and paper, gas and oil, electric power, automotives, aerospace, and telecommunications industries.
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Characteristics of Corrosive Environments Several important characteristics/variables of environments control corrosion behavior: · Relative degree of acidity or alkalinity as measured by pH · Relative degree of oxidizing or reducing conditions measured by the electrochemical potential, E · Temperature of the environment · Presence of detrimental/beneficial species · Velocity/fluid flow rate · Concentration of corrodent in the environment
Each of these characteristics is discussed below. Another key consideration, the effects of biological organisms like bacteria on corrosion, is addressed in a separate section. Acidity/Alkalinity. A wide variety of acid or alkaline conditions are encountered in common environments. As shown in Fig. 1, environments can range from strong acids, represented by low pHs (e.g., pH£1); to neutral environments with pH 7; to strong alkalis, represented by sodium hydroxide and calcium hydroxide at pH 14. The measured pH of several common acids and alkalis are as follows: Solutions
pH
Hydrochloric acid (1N) Sulfuric acid (1N) Boric acid (0.1N) Sodium bicarbonate (0.1N) Ammonium hydroxide (1N) Sodium hydroxide (1N)
0.1 0.3 5.2 11.6 11.6 14
Hydrochloric acid and sulfuric acid are strong acids and have low pH values. Boric acid and sodium bicarbonate are representative of a weak acid and a weak alkaline solution, respectively. Sodium hydroxide is a strongly alkaline solution and results in high pH. The acidity or alkalinity of an environment in combination with its oxidizing/reducing characteristics is critical to the resulting corrosion behavior of various materials. 1
7
Strong acid Hydrochloric, sulfuric, nitric
Fig. 1
Weak acid Boric, citric, phosphoric
Neutral Seawater, tap water
14 Weak alkaline Sodium, bicarbonate, ammonium hydroxide
The pH of several common environments
Strong alkaline Sodium hydroxide, calcium hydroxide
Types of Corrosive Environments
195
Oxidizing
Concentrated nitric acid Wet chlorine
Sulfuric acid Hydrochloric acid (aerated)
Reducing
Fig. 2
Hydrochloric acid (deaerated)
Environments with a wide range of oxidizing/reducing behavior
Oxidizing/Reducing Characteristics. Environments can be characterized by their oxidizing power. An environment that is strongly oxidizing has a great tendency to oxidize materials in contact with it, while an environment that is strongly reducing has a much lesser tendency to oxidize materials. Figure 2 illustrates the oxidizing/reducing power of several common environments. Concentrated nitric acid and wet chlorine environments are strongly oxidizing and promote the oxidation of all but the most noble metals. The oxidizing power of sulfuric acid is a function of concentration and ranges from moderately reducing for dilute acids to more strongly oxidizing for concentrated acids. Deaerated hydrochloric acid is a strongly reducing environment, and the addition of oxidizing species such as dissolved oxygen or ferric (Fe3+) ions will significantly increase oxidizing power. Figure 3 shows oxidizing power as a function of potential, E, and the degree of acidity or alkalinity as measured by pH. Typical ranges of oxidizing power and acidity for several environments are indicated. As stated earlier, nitric acid is a strongly oxidizing acid characterized by low pH values and high potential values. Hydrochloric acid is a reducing
+1.6
E (SHE), V
Nitric acid
Sulfuric acid Atmospheric, seawater, fresh water Boric, Bicarbonate, citric ammonium acid hydroxide
–1.6 Hydrochloric acid
2
Fig. 3
Sodium hydroxide
7 pH
12
Approximate locations of several common environments on a potentialpH diagram. SHE, standard hydrogen electrode
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Corrosion: Understanding the Basics
acid characterized by low pHs and lower potentials. Boric acid is a weaker acid and has a pH representative of mildly acidic solutions. Atmospheric, seawater, and freshwater environments are typically nearly neutral and range from moderately reducing to moderately oxidizing. Sodium hydroxide is a strongly alkaline environment represented by high pH, and bicarbonate and ammonium hydroxide solutions are weak bases with somewhat lower pH values. Metals in contact with these environments may be immune; that is, no attack is exhibited. They also may be passive, where a protective corrosion product layer provides corrosion resistance to the metal, or active, in which case corrosion rates are substantial. The relative locations of immune, passive, and active regions depend very much on the specific metal. Gold is immune over almost the entire E-pH diagram, as it is attacked by only strongly oxidizing environments. Magnesium is a very reactive metal and shows an active region over the bulk of the E-pH diagram. Titanium exhibits a strong tendency to passivate, and large portions of the diagram are covered by the passive region. A number of E-pH diagrams are presented and discussed in Chapter 3. Temperature Effects. Corrosion is an activation-controlled chemical reaction, the rate of which is greatly affected by temperature. Typically, corrosion rate increases significantly as temperature increases. A rule of thumb is that the corrosion rate will double for each 10 °C (18 °F) increase in temperature. The increased corrosion rate results from increased activation energy for chemical and electrochemical reactions, increased diffusion rates in the electrolyte, and increased transport through the electrolyte or environment and across films that may be formed on the metal surface. There are observed reversals in the trends regarding temperature and corrosion rate. These are typically related to the stability of passive films that may form or break down over a fairly narrow temperature range and also can be related to the solubility of corrosive or beneficial gases and species in the environment. An important consideration with respect to temperature is the temperature of the metal surface in contact with the environment. Often the surface temperature of a heating element or heat-exchanger component can be significantly greater than the bulk environment temperature. This can result in significantly increased corrosion rates of these surfaces with respect to the corrosion rate of unheated metal surfaces in contact with the same electrolyte. Detrimental/Beneficial Species. The presence of individual species in an environment, even at trace levels, can have a major impact on corrosion behavior. Chlorides particularly detrimental to materials that depend on a passive film for their corrosion resistance (e.g., stainless steels). In the presence of chloride ion, the passive film can break down locally and result in pitting, concentration-cell corrosion, and in some cases stress-corrosion cracking (SCC). The presence of ammonia in an
Types of Corrosive Environments
environment can greatly increase the corrosion of some alloys. The ammonia ion forms soluble complexes with copper in the metal and results in the breakdown of passive surface films on copper alloys and more rapid corrosion. An environment can also contain beneficial species in the form of corrosion inhibitors that are intentionally added to the environment or in the form of naturally occurring species that help to control corrosion. Beneficial species can promote the development of protective films on the metal surface, or they can interact with detrimental species to reduce the corrosion tendency for the specific metal/environment combination. More detailed information regarding the use of inhibitors for corrosion control can be found in Chapter 9. Velocity/Fluid Flow Rate. The influence of fluid flow rate, or fluid velocity, is a complex variable, and its influence on corrosion behavior is dependent on the alloy, fluid constituents, fluid physical properties, geometry, and corrosion mechanism. The presence of fluid flow can sometimes be beneficial in preventing or decreasing localized attack. For example, type 316 stainless steel has been shown to pit in quiescent seawater but not in moving seawater. When the seawater is moving, the mass transfer rate of oxygen is high enough to maintain a completely passive surface, but in the absence of flow, the mass transfer of oxygen is too slow and the surface cannot remain passive. Under other circumstances, fluid flow can cause a type of erosion of a surface through the mechanical force of the fluid itself. This common process is called impingement, and an example of such erosion is where fluid is forced to turn its flow direction at pipe bends. When solids are present in the liquid, they can cause erosion-corrosion. Both impingement and erosion-corrosion are discussed in Chapter 4. Effect of Concentration. The effect on corrosion rate of increasing or decreasing the concentration of corrodent in the environment to which a metal part is exposed does not follow a uniform pattern because of (a) ionization effects in aqueous solutions and the effects of even trace amounts of water in nonaqueous environments and (b) changes that occur in the characteristics of any film of corrosion products that may be present on the surface of the metal. Typical patterns of the types of corrodent-concentration effects on corrosion rate that may be encountered are illustrated in Fig. 4, which plots the variation of corrosion rate of iron as a function of the concentration of three common inorganic acids in aqueous solutions at room temperature. The rate of corrosion of a given metal usually increases as the concentration of the corrodent increases, as shown in Fig. 4(a) for the corrosion of iron in hydrochloric acid. However, corrosion rate does not always increase with concentration, as shown in Fig. 4(b) and (c) for iron in sulfuric acid and in nitric acid respectively.
197
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Because corrosion is electrochemical and involves anodic and cathodic reactions, process variables influence corrosion rate if they influence one or both reactions. For example, the main cathodic reaction for iron corroding in dilute inorganic acids is 2H+ + 2e– ® H2. The more hydrogen ions available, the faster the rate of the cathodic reaction. In turn, this permits a high rate of anodic dissolution: (Fe ® Fe2+ + 2e–). This is what happens throughout the concentration range in hydrochloric acid solutions. In both nitric acid and sulfuric acid solutions, the hydrogen ion concentration increases with acid concentration. At the higher levels of acid concentration, however, it decreases again. Iron may be used to handle concentrated sulfuric acid at ambient temperatures. Care must be taken to avoid any combination with water, because this dilutes the acid and increases the rate of attack. Also, impurities in the acid or the iron can increase the rate of attack. The corrosion rate of iron (and steel) at room temperature in nitric acid decreases with increasing concentration above approximately 35% because of the formation of a passive oxide film on the metal surface. However, this passive condition is not completely stable. The rate of attack for concentrations of 70% or higher, although low compared to the maximum rate shown in Fig. 4, is still greater than 1.3 mm/year (50 mils/year), making iron and steel unsuitable for use in shipping and storing nitric acids at any concentration. Nitric acid in bulk is usually stored and shipped in type 304 stainless steel, aluminum alloy 3003, or commercially pure titanium (grade 2).
Fig. 4
Effect of acid concentration on the corrosion rate of iron completely immersed in aqueous solutions of three inorganic acids at room temperature. It should be noted that the scales for corrosion rate are not the same for all three charts. As discussed in text, the corrosion rate of iron (and steel) in nitric acid in concentrations of 70% or higher, although low compared to the maximum rate, is sufficient to make it unsafe to ship or store nitric acid in these metals.
Types of Corrosive Environments
199
Metals that have passivity effects, such as Monel 400 in hydrochloric acid solutions and lead in sulfuric acid solutions, corrode at an extremely low rate at low acid concentration at room temperature but lose their passivity at a certain limiting acid concentration. Above the limiting acid concentration, the corrosion rate increases rapidly with increasing acid concentration.
Biologically Influenced Corrosion Biological organisms are present in virtually all natural aqueous environments. In seawater environments, such as tidal bays, estuaries, harbors, and coastal and open ocean seawaters, a great variety of organisms are present. Some of these are large enough to observe with the naked eye (e.g., barnacles, mussels, and clams), while others are microscopic (hence, the term microbiologically influenced corrosion, or MIC). In freshwater environments, both natural and industrial, the large organisms are missing, but there are still many microorganisms, such as bacteria and algae. MIC has been documented for metals exposed to process chemicals, foods, soils, aircraft fuels, human plasma, and sewage, in addition to seawater and freshwater. In all these environments, organisms in the water tend to attach to and grow on the surface of structural materials, resulting in the formation of a biological film, or biofilm. The film itself can range from a
Fig. 5
A gelatinous biofouling slime layer on a heat exchanger tube sheet. The slime layer may be colored by dirt and other debris that accumulates in the gooey mass. Source: Nalco Chemical Company
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Corrosion: Understanding the Basics
Fig. 6
Barnacles attached to the periphery of a high-strength steel rudder, which had originally been coated with an antifouling paint. During use, the paint around the edges had been removed by mechanical action, thus allowing the attachment of barnacles. Partial coverage of such macroorganisms can lead to localized corrosion. Complete coverage can sometimes provide a barrier film and limit corrosion.
microbiological slime film on freshwater heat transfer surfaces (Fig. 5) to a heavy encrustation of hard-shelled fouling organisms on structures in coastal seawater (Fig. 6). The biofilms that form on the surface of virtually all structural metals and alloys immersed in aqueous environments have the capability to influence the corrosion of those metals and alloys. This influence derives from the ability of the organisms to change environmental variables including oxidizing power, temperature, velocity, and concentration. Thus, the value of a given parameter at the metal/water interface under the biofilm can be quite different from that in the bulk electrolyte away from the interface. The result can be the initiation of corrosion under conditions in which there would be none in the absence of the film, a change in the mode of corrosion (that is, from uniform to localized), or an increase or decrease in the corrosion rate.
Industries and Organisms Involved The various industries that have been affected by MIC problems are listed in Table 1. Table 2 lists organisms that cause corrosion in environments
Types of Corrosive Environments
Table 1
201
Industries known to be affected by microbiological corrosion
Industry
Problem areas
Chemical-processing industries
Nuclear power generation
Onshore and offshore oil and gas industries
Underground pipeline industry Water treatment industry Sewage handling and treatment industry Highway maintenance industry Aviation industry Metal working industry
Table 2
Stainless steel tanks, pipelines and flanged joints, particularly in welded areas after hydrotesting with natural river or well waters Carbon and stainless steel piping and tanks; copper-nickel, stainless, brass and aluminum bronze cooling water pipes and tubes, especially during construction, hydrotest, and outage periods Mothballed and waterflood systems; oil and gas handling systems, particularly in those environments soured by sulfate reducing bacteria (SRB), from produced sulfides Water-saturated clay-type soils of near-neutral pH with decaying organic matter and a source of SRB Heat exchangers and piping Concrete and reinforced concrete structures Culvert piping Aluminum integral wing tanks and fuel storage tanks Increased wear from breakdown of machining oils and emulsions
Microorganisms most commonly implicated in biological corrosion pH range
Temperature range in °C
Oxygen requirement
Desulfovibrio (best known: D. desulfuricans)
4–8
10–40
Anaerobic
Iron and steel, stainless steels, aluminum, zinc, copper alloys
Desulfotomaculum (best known: D. nigrificans, also known as Clostridium) Desulfomonas Thiobacillus thiooxidans
6–8
10–40 (some 45–75)
Anaerobic
Iron and steel, stainless steels
… 0.5–8
10–40 10–40
Anaerobic Aerobic
Iron and steel Iron and steel, copper alloys, concrete
Thiobacillus ferrooxidans
1–7
10–40
Aerobic
Iron and steel
Gallionella
7–10
20–40
Aerobic
Iron and steels, stainless steels
Sphaerotilus
7–10
20–40
Aerobic
Iron and steel, stainless steels
S. natans
…
…
…
Pseudomonas
4–9
20–40
Aerobic
P. aeruginosa
4–8
20–40
Aerobic
3–7
10–45 (best at 30–35)
…
Genus or species
Metals affected
Action
Bacteria
Aluminum alloys
Iron and steel, stainless steels Aluminum alloys
Utilize hydrogen in reducing SO42- to S2– and H2S; promote formation of sulfide films Reduce SO42- to S2– and H2S; (spore formers) Reduce SO42- to S2– and H2S Oxidizes sulfur and sulfides to form H2SO4; damages protective coatings Oxidizes ferrous (Fe2+) to ferric (Fe3+) Oxidizes ferrous (and manganous) to ferric (and manganic); promotes tubercle formation Oxidizes ferrous (and manganous) to ferric (and manganic); promotes tubercle formation Oxidizes ferrous (and manganous) to ferric (and manganic); promotes tubercle formation Some strains can reduce Fe3+ to Fe2+ Some strains can reduce Fe3+ to Fe2+
Fungi Cladosporium resinae
Aluminum alloys
in which there would be none without them, accelerate corrosion, or change a relatively slow rate of general corrosion into one with rapid localized penetration of the metal.
Produces organic acids in metabolizing certain fuel constituents
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Sulfate-reducing bacteria (SRB) are the organisms most often cited as the causative agents of MIC for several reasons. They are ubiquitous, easy to culture, and in anaerobic, sulfate-containing environments, they produce hydrogen sulfide (H2S) that can react with iron and steel, stainless steels, copper, aluminum, and zinc. Desulfovibrio, Desulfomonas, and Desulfotomaculum are the three genera of SRB. Table 2 lists the characteristics of these SRB. Sulfate-reducing bacteria frequently cause intense localized attack. As shown in Fig. 7, the resulting pits are hemispherical and often feature smooth interiors. Pits also tend to cluster together, overlapping to form irregularly dimpled surfaces. Acid Producers. Many bacteria produce acids. Acids may be organic or inorganic depending on the specific bacterium. In either case, the acids produced lower the pH, usually accelerating attack. Thiobacillus thiooxidans is an aerobic organism that oxidizes various sulfur-containing compounds to form sulfuric acid (H2SO4). As will be described later in this chapter, these bacteria are sometimes found near the tops of tubercles. Table 2 lists the characteristics of Thiobacillus thiooxidans. Iron/manganese bacteria oxidize ferrous, Fe2+, iron to ferric, Fe3+, iron. Ferric hydroxide, Fe(OH)3, is the result. Some bacteria oxidize manganese (manganous to manganic ions) and other metals. Aerobic Gallionella bacteria, in particular, have been associated with the accu-
Fig. 7
Severely pitted aluminum heat exchanger tube. Pits were caused by sulfate-reducing bacteria beneath a slime layer. The edge of the slime layer is just visible as a ragged border between the light-colored aluminum and the darker, uncoated metal below. Source: Nalco Chemical Company
Types of Corrosive Environments
mulation of iron oxides in tubercles. Table 2 lists the characteristics of Gallionella bacteria. Slime formers are a diverse group of aerobic bacteria. Slime layers (Fig. 5) are a mixture of bacterial secretions called extracellular polymers, other metabolic products, bacteria, gases, detritus, and water. Slime layers contribute to corrosion by forming differential oxygen cells, which lead to crevice corrosion. In addition, slime layers can be efficient scrubbers of oxygen, thus preventing oxygen from reaching the surface. This creates an ideal site for anaerobic SRB growth.
Tuberculation As described in Chapter 4, tubercles are discreet hemispherical mounds over pits on steel surfaces. The process of biologically influenced tubercle formation is a complex one. A number of the reactions that can take place are illustrated for a ferrous alloy in Fig. 8. The volcanolike structure often starts with a deposit of slime-forming and iron-oxidizing bacteria at a point of low flow velocity. This creates an oxygen concentration cell, thus promoting dissolution of iron as Fe2+ under the deposit. As the Fe2+ ions move outward, they are oxidized to Fe3+; this occurs electrochemically as they encounter higher oxygen concentrations and/or by the action of iron bacteria. The resulting corrosion product, Fe(OH)3, mingles with the biodeposit to form the wall of the growing tubercle. When bacteria are present, the tubercle structure is usually less brittle and less easily removed from the metal surface than when bacteria are absent. The outside of the tubercle becomes cathodic, while the metal surface inside becomes highly anodic.
Fig. 8
Schematic diagram of electrochemical and microbial processes involved in tuberculation. Not all of these processes may be active in any given situation. Cl–, chlorides.
203
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Corrosion: Understanding the Basics
As the tubercle matures, some of the biomass may start to decompose, providing a source of sulfates for SRB to use in producing H2S in the anaerobic interior solution. In some cases, the sulfur-oxidizing bacteria may assist in the formation of the sulfates. Depending on the ions available in the water, the tubercle structure may contain some FeCO3, and when SRB are present, some FeS. Finally, if there is a source of chlorides and if the iron-oxidizing bacteria Gallionella are present, a highly acidic, ferric chloride solution may form inside the tubercle. Generally, not all of the above reactions will take place in any single environment. As the individual tubercles on a surface grow under the influence of any combination of reactions, they eventually combine to form a mass that severely limits flow (or even closes it off altogether), leaving a severely pitted surface underneath.
Prevention of MIC Effective MIC prevention requires identification of the bacteria and verification that the bacteria are a factor in corrosion. Microbiologically influenced corrosion should be considered whenever extensive, repetitive pitting corrosion continues, even with an upgraded material. Preventative methods include these: · Providing corrosion monitoring for the presence of MIC causing bacteria. Monitoring should include water and biofilm sampling to determine bacterial counts and biofilm development on surfaces. · Providing periodic cleaning to ensure unimpeded fluid flow and to prevent crevice corrosion · Applying biocide treatments in closed environment systems. Common biocides include chlorine, hydrogen peroxide, and ozone (see Chapter 9 for additional information on biocides). · Using antifouling paints that contain toxic substances, usually copper compounds. They function by slowly releasing copper ions into the aqueous environment, poisoning the mussels, barnacles, and other creatures. · Replacing the iron and steel with more corrosion-resistant materials, such as fiberglass polyvinyl chloride (PVC), polyethylene, and concrete · Creating a nonaggressive environment around the steel by backfilling with gravel or clay-free sand to encourage good drainage (i.e., oxygenating to suppress SRB), making the environment alkaline · Using cathodic protection, although potentials of –0.95 V (or even more negative) versus Cu/CuSO4 are often required. At these potentials, the risk of hydrogen cracking or blistering should be assessed. · Using various barrier coatings, some with corrosion inhibitors and/or biocides
Types of Corrosive Environments
205
Atmospheric Corrosion Atmospheric corrosion is defined as the corrosion or natural degradation of material exposed to the air and its pollutants rather than immersed in a liquid. The rate or degree of degradation varies for different materials and is influenced by several environmental factors. Many of these factors are natural in origin, but some result from man-made sources. Among the latter sources, which are known to affect atmospheric degradation of materials, are the SOx and NOx compounds produced as fossil fuel combustion by-products. These species can react in the atmosphere and result in acid deposition. Although considerable public attention is focused on the effects of acid deposition on our ecosystem, the potential damage to materials may represent the largest economic impact of acid deposition. In fact, it has been reported that atmospheric corrosion accounts for more failures in terms of cost and tonnage than any other single environment. Types of Atmospheres. Atmospheres are often broadly classified as being: · Rural · Industrial · Marine
An atmosphere classified as rural is normally one that does not contain chemical pollutants but does contain organic and inorganic dusts. Its principal corrodent is moisture and, of course, oxygen and carbon dioxide. Arid or tropical atmospheres are special cases of the rural environment because of their extreme relative humidities and condensation. The rural atmosphere is generally the least corrosive. An industrial atmosphere is typically identified as one with heavy industrial manufacturing facilities. These atmospheres can contain concentrations of sulfur dioxide, chlorides, phosphates, nitrates, or other specific industrial emissions. These emissions combine with precipitation or dew to form the liquid corrosive. A marine atmosphere is laden with fine particles of sea salt carried by the winds and deposited on materials. The marine atmosphere is usually one of the more corrosive atmospheric environments. It has been shown that the amount of salt (chlorides) in the marine environment decreases with increasing distance from the ocean and is greatly influenced by wind direction and velocity. Such classifications are, of course, gross oversimplifications of the situation. It is easy to list locations along the seacoast that have heavy industrial pollution in the atmosphere. Such locations are both marine and industrial. Furthermore, two decidedly rural environments can
206
Corrosion: Understanding the Basics
differ widely in average yearly temperature and rainfall and can therefore have considerably different corrosive tendencies. Industrial expansion into formerly rural areas can easily change the aggressiveness of a particular location. Finally, longterm trends in the environment, such as changes in rainfall patterns, mean temperature, and perhaps acid rain, can make extrapolations from past behavior less reliable. Other factors that limit the usefulness of atmospheric-exposure data are the general nonlinearity of weight loss due to corrosion with time and the fact that most atmospheric-corrosion data are presented as an average over the entire test panel surface. Most atmospheric-exposure Table 3 Comparative rankings of 45 locations based on loss of weight (grams) of 10 by 15 cm (4 by 6 in.) specimens of steel Ranking to State College, PA
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45
Location
Norman Wells, N.W.T., Canada Phoenix, AZ Saskatoon, Sask., Canada Esquimalt, Vancouver Island, Canada Detroit, MI Fort Amidor Pier, Panama, C.Z. Morenci, MI Ottawa, Ont., Canada Potter County, PA Waterbury, CT State College, PA Montreal, P.Q., Canada Melbourne, Australia Halifax (York Redoubt), N.S. Durham, NH Middletown, OH Pittsburgh, PA Columbus, OH South Bend, PA Trial, B.C. Canada Bethlehem, PA Cleveland, OH Miraflores, Panama, C.Z. London (Battersea), England Monroeville, PA Newark, NJ Manila, Philippine Islands Limon Bay, Panama, C.Z. Bayonne, NJ East Chicago, IN Cape Kennedy, 0.8 km (0.5 mile) from ocean Brazos River, TX Pilsey Island, England London (Stratford), England Halifax (Federal Building), N.S. Cape Kennedy, 55 m (60 yd) from ocean, 60 ft elevation Kure Beach, NC, 250 m (800 ft) lot Cape Kennedy, 55 m (60 yd) from ocean, 9 m (30 ft) elevation Daytona Beach, FL Widness, England Cape Kennedy, 55 m (60 yd) from ocean, ground level Dungeness, England Point Reyes, CA Kure Beach, NC, 25 mm (80 ft) lot Galeta Point Beach, Panama, C.Z.
Two-year exposure, grams lost
0.73 2.23 2.77 6.50 7.03 7.10 9.53 9.60 10.00 11.00 11.17 11.44 12.70 12.97 13.30 14.00 14.90 16.00 16.20 16.90 18.3 19.0 20.9 23.0 23.8 24.7 26.2 30.3 37.7 41.1 42.0 45.5 50.0 54.3 55.3 64.0 71.0 80.2 144.0 174.0 215.0 238.0 244.0 260.0 336.0
Types of Corrosive Environments
207
data for steels show a decrease in the rate of attack with time of exposure so that extrapolations of such data to times longer than those covered by the exposure data can lead to an overdesign in cross section. Finally, in many cases, the average weight loss per unit area is of less concern than the time to perforation. This factor is more related to localized attack, which can be masked by the averaging of data, as is done in weight-loss determinations. The effect of various environments on the corrosion rate is indicated in Table 3. Table 3 is a compilation of weight-loss measurements for cold-rolled steels after two years of exposure. The most startling feature of Table 3 is the extreme range of corrosion rates: Galeta Point Beach, Panama, is more than 450 times as aggressive as the site at Norman Wells, N.W.T., Canada. This difference in corrosion rate is easily greater than any effect that can be produced by small changes in composition of the steel. This again underscores the fact that in dealing with the corrosion of carbon steels, the alteration of design or environmental factors is usually more effective than changing the grade of steel. Further examination of Table 3 shows that the marine environments tend to be near the aggressive end of the list and that cold environments are generally less aggressive than warm sites. The average yearly temperature cannot, in general, be isolated from the moisture effect, because most of the more tropical exposure sites are also in regions with high humidity. One exception is arid Phoenix, AZ. Atmospheric Corrosion of Various Metals. The average atmospheric corrosion rates of various metals for 10 year and 20 year exposure times are presented in Table 4. Corrosion rates are given for several metals at urban industrial, marine, and rural locations. On a tonnage basis, steel is the most common metal exposed to atmospheric corrosion. Other important metals include aluminum, copper, lead, tin, nickel, and zinc. Aluminum has high corrosion resistance in rural environments Table 4
Average atmospheric-corrosion rates of various metals for 10 year and 20 year exposure times
Corrosion rates are given in mils/year (1 mil/year = 0.025 mm/year). Values cited are one-half reduction of specimen thickness. New York, NY (urban-industrial) Metal
Aluminum Copper Lead Tin Nickel 65% Ni, 32% Cu, 2% Fe, 1% Mn (Monel) Zinc (99.9%) Zinc (99.0%) 0.2% C Steel(a) (0.02% P, 0.05% S, 0.05% Cu, 0.02% Ni, 0.02% Cr) Low-alloy steel(a) (0.1% C, 0.2% P, 0.04% S, 0.03% Ni, 1.1% Cr, 0.4% Cu) (a) Kearney, NJ (near New York City)
Atmosphere La Jolla, CA (marine) Years 10 20
State College, PA (rural)
10
20
10
20
0.032 0.047 0.017 0.047 0.128 0.053 0.202 0.193 0.48
0.029 0.054 0.015 0.052 0.144 0.062 0.226 0.218 …
0.028 0.052 0.016 0.091 0.004 0.007 0.063 0.069 …
0.025 0.050 0.021 0.112 0.006 0.006 0.069 0.068 …
0.001 0.023 0.019 0.018 0.006 0.005 0.034 0.042 …
0.003 0.017 0.013 … 0.009 0.007 0.044 0.043 …
0.09
…
…
…
…
…
208
Corrosion: Understanding the Basics
and lower resistance in industrial and marine environments. Nickel and nickel alloys have high corrosion resistance in marine and rural environments but can exhibit attack in industrial environments containing higher levels of sulfur compounds. Zinc also exhibits its highest corrosion rates in industrial environments, with somewhat lower rates in marine environments and still lower rates in rural environments. Table 4 also provides a comparison between the 10 year corrosion rates of a carbon steel and a low-alloy steel containing alloying additions of phosphorus, nickel, chromium, and copper for increased atmospheric corrosion resistance. Such steels, referred to as weathering steels, are discussed in Chapter 6. Effects of Moisture and Atmospheric Pollutants. Because atmospheric corrosion is an electrolytic process, the presence of an electrolyte is required. This should not be taken to mean that the steel surface must be awash with water; a very thin, adsorbed film of water is all that is required. During an actual exposure, the metal spends some portion of the time awash with water because of rain or splashing and a portion of the time covered with a thin adsorbed water film. The portion of time spent covered with the thin water film depends quite strongly on relative humidity at the exposure site (Fig. 9). This fact has led many corrosion scientists to investigate the influence of the time of wetness on the corrosion rate. These studies have shown that the time of wetness, although an important factor, cannot be considered in isolation when estimating corrosion rates. An excellent example of this fact is demonstrated in Fig. 10, in which the weight loss of iron is plotted as a function of relative humidity for an exposure of 55 days in an atmosphere containing 0.01% sulfur dioxide (SO2). In the lower righthand corner of Fig. 10 is the measured corrosion rate of iron exposed for the same time in an
Fig. 9
Relationship between the corrosion rate of iron and the relative humidity in an environment containing 0.01% SO2. Exposure period: 55 days. The corrosion rate of iron exposed for 55 days at 100% relative humidity is shown for comparison.
Types of Corrosive Environments
SO2-free atmosphere at 99% relative humidity. The increase in corrosion rate produced by the addition of SO2 to the atmosphere is substantial. Another feature of interest is the apparent existence of a critical humidity level below which the corrosion rate is small. As shown in Fig. 10, the critical humidity in an SO2-containing atmosphere is approximately 60%. To better appreciate the effects of pollutants on the atmospheric corrosion of metals, one may consider once again the environmental aspects of automotive corrosion. Before 1950, corrosion of externally exposed metal on automobiles was not considered a serious problem. By 1955, initial signs of corrosion had surfaced. During the 1960s, type 430 stainless steel containing 17% Cr was failing due to pitting and crevice corrosion. This stainless steel grade was widely used for exterior trim. As a result of localized corrosion problems, a modified 434 grade with 17% Cr and 1% Mo was developed. Severe galvanic corrosion problems with painted auto body steel surfaced during the 1960s, especially in trim areas where solid stainless steel was attached to the auto body. By 1970, corrosion perforation of painted auto body steel was a problem which grew to major importance throughout the 1970s. The susceptibility and high incidence of pitting and blush and bloom (flat finish with milky appearance) of anodized aluminum exterior trim was recognized by 1980 and by 1985, type 434 stainless steel in the same application was exhibiting signs of localized corrosion unless it was properly produced. High susceptibility to cosmetic corrosion, including filiform corrosion, persisted throughout the 1990s. Although some of these problems can be attributed to poor design and maintenance, there is also a direct correlation with atmospheric
Fig. 10
Effect of relative humidity and atmospheric pollution (environment containing 0.01% SO2) on the rusting of iron. A critical humidity near 60% relative humidity is observed.
209
SO2 emissions
8 6
20
NOx emissions 10 Salt usage
4 2 0
Emissions, millions of tons
Corrosion: Understanding the Basics
Road salt usage, millions of tons
210
0
1945 1950 1955 1960 1965 1970 1975 1980 1985 1990
Year
Fig. 11
Annual emissions and road salt usage in the United States
conditions. Figure 11 plots the annual emissions output and road (deicing) salt usage in the United States from 1945 through 1990. As the use of road salts (sodium chloride and calcium chloride) increased and levels of atmospheric pollutants (emissions) continued to rise, the corrosion problems summarized in the previous paragraph began to manifest themselves. Chlorides are known to be aggressive to metals and atmospheric pollutants such as SO 2 , NO x , H 2 S, and NH 4 , and particulate matter are known to have a negative impact on corrosion. Take, for example, the reaction involving the SO 2 molecule striking a steel surface and converting to sulfuric acid as a result of hydrolysis and oxidation. This introduces reactive ions, increased conductivity, and a lower pH. The presence of hygroscopic sulfate corrosion products also lowers the critical humidity for the corrosion reaction. Note that the most severe automotive corrosion problems during the 1970s correspond with the highest emission levels and road salt usage shown in Fig. 11. Realization of improved corrosion performance began after the U.S. Clean Air Act of 1977, with public awareness of the damaging effects of road salts and better quality control by automobile manufacturers. Corrosion Control. Atmospheric corrosion is controlled by selecting more corrosion resistant materials. For example, aluminum or stainless steels can be used where the corrosion of steel parts is undesirable. Metallic and organic coatings are commonly used on steel to provide atmospheric corrosion resistance. The metallic coatings can be sacrificial, such as zinc coating on galvanized steel, or they can be a noble metal coating that acts as a barrier. Removal of moisture from the atmosphere to levels below the critical relative humidity will control corrosion. Air conditioning is commonly used to control the relative humidity in the atmosphere for protection of communications equipment and computer components.
Types of Corrosive Environments
211
Underground/Soil Corrosion Soil corrosion may be considered to encompass all corrosion taking place on buried (underground) structures. Underground storage tanks, gas and oil transmission pipelines, distribution pipelines, and many other structures buried in soil are susceptible to underground corrosion. Ordinary steel and cast irons are the two most common metallic materials used for such underground structures. Aluminum, stainless steel, and copper are also used. Nonmetallic materials are being used increasingly for underground applications. Concrete is used as footings, piers, tanks, and piping. Thermoplastics, most notably polyvinyl chloride (PVC), acrylonitrilebutadiene-styrene (ABS), and polyethylene (PE) have replaced metals in many piping applications. Glass fiber reinforced thermosetting plastics are used for the underground storage of a wide range of products (e.g., water and gasoline).
Factors Affecting Underground/Soil Corrosion Important soil characteristics that control the corrosivity of soils to steel and other metals include the following: · Aeration and water-retention (permeability) characteristics of the soil · Dissolved salt content and resistivity of the soil · Soil acidity · Presence of ionic species in the soil—for example, chloride or microbiologically active species, such as bacteria
The resulting corrosion resistance depends on the combination of the corrosivity of the soil and the corrosion behavior of the metal under those conditions. Aeration and water-retention characteristics are the chief physical attributes of soil. Coarse soils (sands and gravels) that have good drainage and ample aeration corrode steel at a rate approaching that of the local atmosphere. Clay and silt soils are characterized by fine texture, high water retention, poor aeration, and poor drainage, all of which increase corrosion rates for steel. The most severe corrosion usually takes place at low elevations in poorly drained soils such as clays and tidal marshes where there is minimal aeration. If these soils are dug up and then backfilled, drying and shrinkage cause them to become partly aerated, which, in turn, creates differential aeration cells. Salts often accumulate in the backfilled areas and lower the soil resistivity. In
212
Corrosion: Understanding the Basics
contrast, corrosion occurs to a lesser degree in well-aerated, welldrained, porous soils where salts are readily washed away by the rain. Corrosion by differential aeration can result from various conditions. For example, when a pipe passes through two soils that differ in oxygen permeability, a galvanic current flows from the more poorly aerated (anodic) surface of the pipe through the soil to the better aerated (cathodic) surface, as shown schematically in Fig. 12. Lower oxygen concentrations commonly occur at the bottom of a buried steel structure, where the soil is more compact and farther from the source of oxygen in the atmosphere. This can give rise to an oxygen concentration cell. The lower-oxygen environment in contact with the steel at the bottom of the pipe can cause this area to become anodic with respect to the steel at the top of the pipe, which is in contact with soil that contains higher levels of oxygen. Corrosion current leaves the steel at the bottom of the pipe, resulting in corrosion at the anodes, and travels through the soil to the top of the pipe, where the cathodic reaction is the reduction of oxygen (Fig. 13). This type of corrosion also can occur if the normal water table lies between the top and bottom of a large- diameter pipe. Oxygen-concentration cells are also produced at random sites where sticks, stones, and other foreign material in the backfill contact the pipe. Soil acidity can also influence underground corrosion. The acidity or alkalinity of most soils in the United States is stable because of the buffering action of the soluble minerals available. However, soils across the country differ in pH from less than 4.5 (very acidic) to greater than 9.5 (strongly alkaline). The corrosion response of a metal in contact with soil varies widely, depending on acidity. Steel, for example, is corroded in acidic environments but remains passive in moderately alkaline environment, whereas aluminum is passive in a neutral environment but is corroded in either extremely acidic or strongly alkaline environments.
Fig. 12
Corrosion cell on a metal surface buried in soil. The corrosion current leaves the metal surface at an anode and enters the soil. The current flows from the anode to the cathode through the soil by ionic conductivity and enters the metallic structure at the cathode. The current flows from the cathode through the metallic structure to the anode by electrical conductivity and completes the corrosion cell. In most cases in naturally occurring soils, that portion of the pipe lying in the more conductive (poorly aerated) soil is the anode; that in the less conductive (well aerated) soil is the cathode
Types of Corrosive Environments
213
Dissolved Salts and Soil Resistivity. The most corrosive soils are those that contain large concentrations of soluble salts. Because of the presence of salts, such soils have relatively high electrical conductivites (or low electrical resistivities). The least corrosive soils have low concentrations of soluble salts and high resistivities. Resistivity measurements can be obtained readily and provide as much information on corrosion characteristics as measurements of any other single soil property. Consequently, resistivity is the property most often used to approximate the aggressiveness of a soil. Observations of soil drainage and/or measurements of pH supplement resistivity measurements. The following table lists the general relationship that exists between soil resistivity and corrosion of ferrous metals; however, because of other factors, this relationship may not always be valid. Soil resistivity, W · m
<7 7–20 20–50 >50
Classification
Very corrosive Corrosive Moderately corrosive Mildly corrosive to noncorrosive
Presence of Biological Species. It has been estimated that 50% of all failures of buried metal structures are due to MIC. The primary culprits are SRB and acid-producing bacteria such as Thiobacillus (refer to the earlier section “Biologically Influenced Corrosion” in this Chapter). In view of the effects of bacteria, the best means for modifying a natural environment is to displace organic soil, with its bacterial population and nutrients, from around the buried structure and to surround the structure with mineral backfill (gravel or clay-free sand) that provides good drainage and adequate aeration. The backfill can be treated with biocides to provide an environment that inhibits growth and reproduction.
Types of Underground/Soil Corrosion Galvanic corrosion is a concern in soils where dissimilar materials are in contact. A classic example is the use of copper alloy (brass or Higher-oxygen cathode
Lower-oxygen anode
Fig. 13
An oxygen concentration cell on steel in soil
214
Corrosion: Understanding the Basics
bronze) valves and fixtures with steel piping. The copper alloy component becomes a cathode and accelerates the corrosion of steel. Often the accelerated attack is most severe for the steel immediately adjacent to the copper alloy component. Table 5 presents a galvanic series in typical neutral soil. When two dissimilar metals are electrically coupled in contact with moist soil, the more positive member of the couple becomes a cathode and the more negative member becomes an anode. The anodic member is corroded more severely. Graphite and copper alloys coupled to steel accelerate the corrosion of the latter. The more active metals in the series (magnesium, zinc, and, in some cases, aluminum) can be used to cathodically protect steel. Where aluminum is used, it must remain active. Another observation from Table 5 is that the presence of corrosion products (rust) or mill scale on steel make the steel more cathodic. A consequence of this is that a fresh, bare steel coupled to a corroded steel or one covered with mill scale shows accelerated attack. Stray-current corrosion is a special form of corrosion observed in buried structures. It differs from natural corrosion in that it is caused by an externally induced electrical current and is basically independent of environmental factors such as oxygen concentration or pH. Environmental factors may enhance other corrosion mechanisms involved in the total corrosion process, but the stray-current corrosion portion of the mechanism is unaffected. Stray currents are defined as currents that follow paths other than their intended circuit. A current leaves the intended path because of poor electrical connections within the circuit of poor insulation around the intended conductive material. It then passes through soil, water, or any other suitable electrolyte to find a low-resistance path, such as a buried metal pipe or some other metal structure, and flows to and from that structure, causing accelerated corrosion. Sources of stray currents include equipment using the earth as a ground return, such as electric Table 5
Practical galvanic series of metals and alloys in neutral soils and water
Metal or alloy(a)
Carbon, graphite, coke Mill scale on steel High-silicon cast iron Copper, brass, bronze Low-carbon steel in concrete Lead Cast iron (not graphitized) Low-carbon steel (rusted) Low-carbon steel (clean and shiny) Commercially pure aluminum Aluminum alloy (5% Zn) Zinc Magnesium alloy (Mg-6Al-3Zn-0.15Mn) Commercially pure magnesium
Potential(b), V
+0.3 –0.2 –0.2 –0.2 –0.2 –0.5 –0.5 –0.2 to –0.5 –0.5 to –0.8 –0.8 –1.05 –1.1 –1.16 –1.75
(a) Arranged from most noble to most active. (b) Typical potential normally observed in neutral soils and water measured with respect to copper sulfate reference electrode
Types of Corrosive Environments
Single structure
Fig. 14
215
Multiple structures
Two examples of stray-current corrosion on buried structures. Black areas indicate anodic areas (i.e., current leaves structure).
railways or welding generators, as well as fixed installations that have inadvertently allowed current to escape, for example, cranes or direct current power transmission lines. Two instances of stray-current corrosion are shown schematically in Fig. 14. Stray-current corrosion results from currents flowing through the soil that enter a buried metallic structure in one area, then leave the metallic structure, and then reenter the soil at another location. The area where the current reenters the soil becomes anodic, and rapid corrosion can occur. For the single structure shown in Fig. 14, the current flowing through the soil enters the structure on the left side, flows through the metal structure, and reenters the environment on the right side. Severe attack is observed at the area where the current reenters the environment. For the case of multiple structures shown schematically in Fig. 14, the current flowing through the upper pipeline by electrical conductivity leaves that pipeline and jumps onto the lower pipeline, where it is electrically conducted at the end of that pipeline and reenters the soil. Severe attack is observed on both pipelines in the area where the current leaves the structure. The corrosion of the upper pipeline could be eliminated by attaching an electrically conducting bonding strap that would tie the two metallic structures together. In that case, the current flowing along the upper pipeline could go through the metallic conductor to the lower pipeline without reentering the soil and, therefore, would not affect the corrosion behavior of the pipeline in the soil.
Corrosion Control The most economical and effective method of combating corrosion of steels underground has been cathodic protection. Both applied current from rectifiers and current from sacrificial anodes have been used, either alone or in conjunction with organic coatings. More detailed information on these corrosion control methods can be found in Chapters 9 and 10.
216
Corrosion: Understanding the Basics
Natural and Treated Waters Natural waters exhibit a wide variety of dissolved ionic species— from nearly pure rain water to concentrated brines. The general classifications of natural waters are fresh, brackish, seawater, geothermal, and brines. The important factors controlling the corrosivity of a water include its acidity, oxidizing power, conductivity, and the presence of specific detrimental or beneficial species. The presence of solids in a water can also significantly affect corrosivity, as can flow rate, temperature, and presence of biological agents. Seawater is a highly conductive environment with approximately 3.4% salt concentration. Other factors that contribute to the corrosivity of seawater include oxygen concentration, velocity, temperature, and presence of marine biological species. Figure 15 shows the relative loss of metal thickness for a steel piling after five years of exposure in seawater at Kure Beach, NC. A variety of corrosive conditions exist from the bottom of the piling in the sea to the portion of the piling exposed to marine atmospheric corrosion. In the soil (zone 5) and the continuously submerged area (zone 4), the corrosion rate is relatively low. The corrosion rate increases at the upper regions of the immersed area and the area in the tidal zone (zone 3), except for a portion immediately above the mean low-tide level and the below-the-mean high-tide level. The corrosivity increases again significantly in the splash zone (zone 2) and then drops off again in the marine atmospheric corrosion zone (zone 1). The tidal and splash zones are considered to be the most corrosive because of the higher oxygen concentration in these areas and the effect of the seawater velocity. The lower corrosion rate of the piling in zone 3 in Corrosion zone
Atmospheric zone 1
Splash zone 2 (above high tide)
Tidal zone 3
Submerged zone 4
Subsoil zone 5
Fig. 15
Description
The portion of the elevated structure subject to a marine atmosphere, including sea mist and high relative humidity, but without significant wetting by splash from waves The portion above the level of mean high tide that is subject to wetting by large droplets of seawater The portion of the structure between mean high tide and mean low tide, alternately immersed in seawater and exposed to a marine atmosphere The portion of the structure from about 0.3 to 1 m (1 to 3 ft) below mean low tide down to the mud line The portion below the mud line, where the structure has been driven into the ocean bottom
Zones of corrosion for steel piling in seawater, and relative loss of metal thickness in each zone
Types of Corrosive Environments
217
Fig. 15 is reported to be due to the buildup of marine life (barnacles, for example) in this area, which provide some corrosion protection if coverage is complete (refer to Fig. 6). In addition, steel at the tidal zone may act cathodically and receive some protection from corrosion just below the tidal zone in the case of a continuous steel pile. Treated waters are classified as potable, demineralized, boiler feed, processed, and cooling. Potable waters show a variety of corrosiveness, much of which is the result of the degree of water hardness. Calcium and magnesium salts in water can precipitate out and form hard scale on pipe surfaces. Corrosion control under these conditions involves striking a balance between sufficient scaling properties that allow a buildup of thin scale for corrosion protection but will not plug and constrict the piping system. If the degree of water hardness is very low, then there is no scaling tendency, and the water can be more corrosive. Demineralized water has soluble salts removed. As the resistivity of the water increases, the corrosion rate of metals typically decreases. For boiler feedwater, oxygen commonly is removed by mechanical or chemical means and the pH adjusted to mildly alkaline conditions. Other inhibitors can also be added. For process water and cooling waters, a variety of chemical treatments are available to control oxygen concentration. The same benefit can be gained by the use of inhibitors. Stagnant water conditions are often more corrosive than flowing conditions. Rapid failures in stainless steel components have resulted from stagnant water, which allows the buildup of deposits with subsequent accelerated attack beneath these deposits.
Understanding Corrosion in Acids Many of the severe corrosion problems encountered in the processing industries involve acids. These environments are naturally corrosive. Metal dissolves at the anodic areas by the oxidation reaction of metal atoms becoming metal ions in solution. The balancing reduction reaction at the cathodes is the consumption of the hydrogen ions in the acid with the evolution of hydrogen gas. The overall reaction is Metal + acids(H+) ® hydrogen(H2) + metal ions (M+) The rate of oxidation or corrosion of the metal is often controlled by the rate of the reduction reaction. In a concentrated acid with a high concentration of hydrogen ions (H+), there is a large capacity for consuming electrons and promoting the corrosion reaction. In the sections that follow, the corrosion behavior of steels in a variety of acids will be described.
218
Corrosion: Understanding the Basics
HCI concentration
The response of steels to these acids is a function of acid concentration and temperature. Hydrochloric, sulfuric, and nitric acids are all strong acids, and yet the corrosion behavior of steel in them differs widely. Two of the keys to understanding this diversity of behavior are the oxidizing power of the acid environment and the tendency for passivation of steel in the environment. It is useful to contrast the corrosion behavior of steel in hydrochloric acid and sulfuric acid as a function of the concentration of the acid. Corrosion Rate Versus Concentration of Acid. The corrosion rate of steel as a function of hydrochloric acid (HCl) concentration is shown in Fig. 4(a) and schematically in Fig. 16. Hydrochloric acid is a strong and highly reducing acid. No passive film is formed in hydrochloric acid and steel corrodes by active dissolution. As the concentration of hydrochloric acid increases, the corrosion rate of the steel increases. The shape of this curve is characteristic of an activation-controlled dissolution of an active metal. The hydrochloric acid concentration is comparable to increasing oxidizing power along the y-axis, and the corrosion rate is comparable to the log of the current for the oxidation reaction along the x-axis. Data for the corrosion of steel by sulfuric acid at room temperature are presented in Fig. 4(b) and Table 6. Very dilute sulfuric acid is only moderately corrosive. As the concentration of sulfuric acid increases from very dilute up to values from 10 to 50%, the corrosion rate increases significantly. At concentrations from 60 to 90%, the corrosion rate drops significantly and remains fairly constant. This behavior is shown schematically in Fig. 17 for the corrosion rate of steel as a function of sulfuric acid concentration. As the sulfuric acid concentration increases, the oxidizing power of the environment increases. At dilute concentrations of acid, this increased oxidizing power is accompanied by an increased corrosion rate of the steel while the steel is actively dissolving. At a critical concentration of acid, the corrosion rate drops
Active dissolution
Corrosion rate
Fig. 16
Corrosion rate of steel as a function of hydrochloric acid concentration
Types of Corrosive Environments
Table 6
219
Corrosion of steel by sulfuric acid, H2SO4, at room temperature Corrosion rate
H2SO4, %
mm/year
0.005 0.05 0.5 1.0 3.0 5.0 10.0 50.0 60.0 70.0 80.0 90.0
mils/year
0.2 0.4 3 5 22 30 58 41 1.5 0.5 0.5 0.5
8 15 120 200 875 1200 2300 1600 60 20 20 20
rapidly to a significantly lower value and remains fairly constant even at increasingly higher levels of concentration. This lower corrosion rate results from the stable passive film that builds up on steel under the more oxidizing conditions. Once the passive film forms, the corrosion rate is independent of oxidizing potential as long as the passive film is stable. Influence of Oxidizing Power and Passivation. The differences in behavior between steel in hydrochloric acid and steel in sulfuric acid can be readily understood in terms of the oxidizing power of the environment and the relative tendency to form passive films. In hydrochloric acid, no passivity of steel is observed, and the corrosion rate continues to increase with increasing oxidizing power of the environment. In sulfuric acid, there is a critical oxidizing power at which passivity becomes stable, and the corrosion rate above this oxidizing power level drops off dramatically. The corrosion resistance of steel under these conditions depends on the tenacity and durability of the passive film. If the film is broken down mechanically or chemically, the corrosion rate of steel will greatly increase in the more concentrated sulfuric acid.
H2SO4 concentration
Passivity
Active dissolution
Corrosion rate
Fig. 17
Corrosion rate of steel as a function of sulfuric acid (H2SO4) concentration
220
Corrosion: Understanding the Basics
Corrosion by Sulfuric Acid Sulfuric acid (H2SO4) is the largest-volume inorganic acid currently in use and is generally considered to be the most important industrial chemical. The principal uses of H2SO4 are for fertilizer manufacture, petroleum refining, production of other chemicals, manufacture of rayon and other textiles, pickling steel and brass, refining metals, manufacture of pigments and colors, explosives, manufacture of coke and coal tar products including dyes and drugs, storage batteries, production of natural and synthetic rubber, and synthetic detergents. Sulfuric acid is made by the contact process, in which elemental sulfur or sulfur-containing waste is burned to form sulfur dioxide (SO2). Sulfur dioxide is converted to sulfur trioxide (SO3) by contact with a vanadium catalyst. Reaction of this gas with water produces sulfuric acid: SO3 + H2O ® H2SO4 Sulfur trioxide is soluble in 100% H2SO4, so if the SO3 (gas) is bubbled through the acid, a “fuming acid” called oleum is produced. Most of the H2SO4 produced is shipped to the customer in three concentrations, approximately 78 and 93%, and fuming (oleum).
Materials Selection Guidelines for Sulfuric Acid The corrosiveness of H2SO4 depends on many factors, particularly temperature and concentration. Strong, hot conditions present the greatest problems, and few materials except platinum, tantalum, fluorocarbon plastic, and brick-lined steel will resist 60 to 98% H2SO4 at 120 °C (250 °F). However, other variables also influence the resistance of materials to H2SO4. The presence of oxidizing or reducing contaminants, velocity effects, solids in suspension, and galvanic effects can alter the serviceability of a particular material of construction. It is unwise to select materials of construction for equipment that will handle H2SO4 solely on the basis of published corrosion data unless the conditions involved are adequately and specifically covered by the reference data. Seemingly minor differences in impurities or environmental conditions can significantly affect actual service corrosion rates. Impurities such as halides generally increase corrosion. Aeration or the presence of oxidizing agents generally accelerates corrosion of nonferrous materials and reduces corrosion of stainless alloys, but the extent of these effects depends on specific conditions. Hot-wall effects are frequently overlooked, and heating coils made of the same material as the containing vessel can corrode rapidly while the condition of the vessel itself remains satisfactory.
Types of Corrosive Environments
221
It is therefore, advisable to consider all general corrosion data only as an indicator of relative resistance and as a guide by which the limiting conditions of materials may be further reviewed. Final selection of materials for specific equipment depends, of course, on such factors as allowable corrosion rate, desired mechanical and physical properties, fabrication requirements, availability, and cost.
Use of Steel in Sulfuric Acid General Resistance. Ordinary carbon steel is widely used for the handling of sulfuric acid in concentrations greater than 70% under static or low-velocity conditions (<0.9 m/s, or 3 ft/s). Storage tanks, pipelines, tank cars, and shipping drums made of steel are very common for 78, 93, and 98% and stronger acids, such as oleum. Pumps and valves may be made of alloy steel because of erosion-corrosion of carbon steel. Much of the equipment if contact sulfuric acid manufacturing plants is made of plain carbon steel. Steel is rapidly attacked by the more dilute sulfuric acids. Figure 18 shows the corrosion of carbon steel by strong sulfuric acid as a function of temperature. Most of the tests were made on ordinary low-carbon steel (0.20% C) with exposure periods of 48 h. Specimens were prepared by abrading the surface with No. 120 emery cloth. The curves in Fig. 18 comprise an isocorrosion diagram and represent corrosion rates of 5, 20, 50, and 200 mils/year (0.13, 0.5, 1.3, and 5 mm/year). These are isocorrosion or constant corrosion lines. In other words, the outlined areas represent regions where corrosion rates of 0 to 5, 5 to 20, 50 to 200, and >200 mils/year would be expected. The corrosion of steel by strong sulfuric acid is complicated because of the 325 Over 200 mils/year 275
Temperature, F
200 225
50–200 50-200 mils/year
175 50 125
20–50 20-50 mils/year
20
5
75 5–20 5-20 mils/year 25 60
65
70
75
80
85
90
0–5 0-5 mils/year 95
100
105
110
Sulfuric acid concentration, %
Fig. 18
Corrosion of steel by sulfuric acid as a function of concentration and temperature
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peculiar dips in the curves and the rapid increase in corrosion in the neighborhood of 100 to 101% acid. The narrowness of this range means that the acids must be carefully analyzed in order to obtain reliable corrosion data. The dips (increased attack) around 85% are more gradual and less difficult to establish. Localized attack can occur even at flow velocities within the prescribed limits (under 0.9 m/s, or 3 ft/s, as mentioned earlier). Discontinuities such as short-radiused elbows, excessive penetration of welds, and pipe mismatch may cause sufficient downstream turbulence to disturb the protective sulfate film, resulting in high corrosion rates. Weldments must be thoroughly inspected to ensure that they contain no slag, surface porosity, laps, excessive penetration, or other welding defects that might initiate accelerated corrosion. In addition, steel vessels and piping should be free of mill scale, or serious pitting or nonuniform corrosion may occur. Hydrogen grooving is another form of localized attack that occurs on vertical or inclined surfaces exposed to the liquid phase. During the corrosion of steel by H2SO4, atomic hydrogen is evolved. If produced in sufficient quantities, the hydrogen combines to form small bubbles that stream along preferred paths on vertical and inclined surfaces, disrupting the soft protective iron sulfate film. Channels and deep grooves may eventually form. Grooving is commonly observed in the tops of horizontal manways on the side of storage tanks and on the top 180° of horizontal pipe runs (Fig. 19). In piping, stagnant conditions promote grooving; therefore, a minimum velocity of 0.3 m/s (1 ft/s) is often recommended. Grooving, combined with erosion-corrosion, also occurs on the sidewalls of tanks (Fig. 20). Location of the liquid inlet in the roof near the shell has, in at least two cases, resulted in combined erosion-corrosion and hydrogen grooving, which caused catastrophic rupture of large storage tanks.
Fig. 19
Hydrogen grooving of a 75 mm (3 in.) diameter steel elbow. The elbow was sectioned; the top half is shown.
Types of Corrosive Environments
Fig. 20
Hydrogen grooving on the sidewall of an H2SO4 storage tank
Use of Cast Irons in Sulfuric Acid Gray cast iron is at least as resistant to corrosion as steel in the 65 to 100% acid concentration range. Gray cast iron is also less sensitive to velocity than steel and is frequently used up to 1.7 m/s (5.6 ft/s). Ductile irons are slightly less resistant than gray irons. Higher silicon contents (3.5% versus the normal 1.8 to 2.8%) can improve corrosion resistance. High-Silicon Cast Iron. Another material that has long been used for handling sulfuric acid is high-silicon iron. Iron with 14.5% Si has exceptional resistance to sulfuric acid in all concentrations to 100% up to the atmospheric boiling points. Corrosion rates are normally less than 0.13 mm/year (5 mils/year).
Use of Stainless Steels in Sulfuric Acid The common stainless steels are not often used for handling straight sulfuric acid. Type 316 stainless steel (18Cr-12Ni with 2% Mo) is sometimes used in dilute acid solutions containing other ingredients. Other stainless steels sometimes used include type 317 and silicon stainless steels containing 5 to 6% Si. The stainless steels show borderline passivity effects; therefore, they are difficult to test, and it is difficult to predict performance or give reliable and reproducible corrosion data. The 20-type alloys are usually the first considered when a sulfuric acid environment is too corrosive for the use of steel, 300-series stainless steels, or cast iron. This group contains both wrought alloy 20Cb-3 (Fe-34Ni-20Cr-3.3Cu-2.5Mo) and cast alloy CN-7M containing 19 to 22% Cr, 27.5 to 30.5% Ni, 2.0 to 3.0% Mo, and 3.0 to 4.0% Cu. These alloys possess excellent (less than 0.13 mm/year, or 5 mils/year) corrosion resistance over the entire concentration range at temperatures up to about 50 °C (125 °F) and possess good resistance (less than 0.5 mm/year, or 20 mils/year) up to about 65 °C (150 °F). Higher temperatures can be used in the more dilute and in the strong acids.
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Use of Nickel Alloys in Sulfuric Acid The corrosion rates of nickel-base alloys, in general, increase with acid concentration up to approximately 90%. Higher concentrations of the acid are generally less corrosive. A broad view of the relative performance of various stainless steels and nickel alloys in pure sulfuric acid is illustrated in Fig. 21. In this graph, the temperatures at which the corrosion rate of an alloy in a given concentration of the acid exceeds 0.5 mm/year (20 mils/year) are plotted as isocorrosion curves. At low acid concentrations, the nickel-chromium-molybdenum alloys show a significantly higher resistance than type 316 stainless steel. Alloy 20, a high-nickel stainless steel, shows similar behavior. These highnickel-chromium-molybdenum alloys can be used only to moderate temperatures in the intermediate and high concentrations of H2SO4. The nickel-molybdenum alloys B and B-2 can be used to higher temperatures for all concentrations of acid. However, in the presence of oxidizing species and in aqueous acid systems, alloys B and B-2 suffer serious corrosion.
Fig. 21 mils/year).
Comparative behavior of several nickel-base alloys in pure sulfuric acid (H2SO4). The isocorrosion lines indicate a corrosion rate of 0.5 mm/year (20
Types of Corrosive Environments
The presence of oxidizing impurities can be beneficial to nickel-chromium-molybdenum alloys shown in Fig. 21 because these impurities can aid in the formation of passive films that retard corrosion. Another important consideration is the presence of chlorides (Cl–). Chlorides generally accelerate the attack; the extent of acceleration differs for various alloys.
Other Metals Used in Sulfuric Acid Zirconium has excellent resistance to sulfuric acid up 50% concentration at temperatures to boiling and above. From 50 to 65% concentration, resistance is generally very good at elevated temperatures, but the passive film is a less effective barrier. Tantalum has resistance to sulfuric acid in all concentrations up to 98% at temperatures as high as 200 °C (390 °F). It is inert to dilute acid even at boiling temperatures and is not attacked by concentrated acids at temperatures below 150 °C (300 °F). Lead resists sulfuric acid but its protective sulfate film is increasingly solubilized above 90% concentration. The film is easily damaged by erosion or abrasion even at low velocities, with the rate of attack increasing with concentration.
Nonmetallic Materials Used in Sulfuric Acid Nonmetallic materials of construction have wide application in H2SO4. Most of these materials have good corrosion resistance to the pure acid, particularly in dilute concentrations, and are primarily restricted by their mechanical properties at temperature. They are relatively unaffected by most inorganic contaminants, except for such strongly oxidizing agents as nitric acid, peroxides, and dichromates. Commonly used materials include these: · Brick linings (used for most severe sulfuric acid conditions) · Polyvinyl chloride (excellent resistance to about 93% concentration at ambient temperatures) · Plastic linings for pipe (see Table 7) · Polyethylene (high degree of resistance up to 98% concentration at ambient temperature) · Fluoroplastics (see Table 8) · Glass-reinforced polyester resins (show virtually no attack in dilute sulfuric acid at 90 °C, or 200 °F, and can be used up to 75% concentration at temperatures less than 65 °C, or 150 °F) · Carbon (used for hot acid at concentrations above 60%) · Graphite (excellent corrosion resistance to all but highly oxidizing concentrations)
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Table 7
Suitability of plastic-lined piping systems in sulfuric acid, H2SO4 Maximum temperature(a) °C °F
Material
Polyvinylidene chloride (PVDC)
Polypropylene (PP)
Polyvinylidene fluoride (PVDF)
Polytetrafluoroethylene (PTFE)
52 24 NR 93 80 65 50 120 105 80 65 260
126 75 NR 200 175 150 120 250 220 175 150 500
Concentration of H2SO4, %
<16 30 >60 <60 93 96 98 <16 30–60 85–93 94–98 0–100
(a) Temperature may be further limited by mechanical conditions under vacuum. NR, not rated
Table 8
Suitability of fluorinated plastic linings in H2SO4
Material
Polytetrafluoroethylene (PTFE) Perfluoroalkoxy (PFA) Fluorinated ethylene propylene (FEP) Ethylene-chlorotrifluoroethylene (ECTFE) Ethylene-tetrafluoroethylene (ETFE) Polyvinylidene fluoride (PVDF)
Maximum temperature °C °F
260 260 205 150 150 120 105 80 65
500 500 400 300 300 250 220 175 150
H2SO4 concentration, %
0–100 0–100 0–100 <98 <98 <16 30–60 85–93 94–98
· Rubber linings, for example, butyl rubber and neoprene (good resistance to 50% concentration at modest temperatures and resistant to 75% acid at ambient temperature) · Glass linings (glass-lined steel widely used for sulfuric acid service)
Corrosion by Nitric Acid Nitric acid (HNO3) is typically produced by the air oxidation of ammonia (NH3). This catalyzed reaction takes place at very high temperatures. The gaseous oxidation product is condensed to an aqueous liquid of about 65% concentration. During the high temperature oxidation, corrosion of the plant materials is of secondary concern. The elevated operating temperatures dictate that the high-temperature properties of the materials are the primary design consideration. Corrosion considerations prevail during and after condensation and at lower temperatures. The concentration of HNO3 up to 99% requires secondary processing to remove excess water. This involves mixing 65% HNO3 with another substance having a greater affinity for water (such as H2SO4) and then separating the mixed acids by distillation and condensation processes.
Types of Corrosive Environments
Fig. 22
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Comparison of corrosion of aluminum alloy 3003 and type 304 stainless steel in HNO3
Materials Selection Guidelines for Nitric Acid Commercially produced HNO3 is available in concentrations from 52 to 99%. Nitric acid over 86% is described as fuming. Nitric acid up to 95% is stored and shipped in type 304 stainless steel. Concentrated acid above 95% is handled in Aluminum Association (AA) aluminum alloys 1100 or 3003. Figure 22 shows the reason for this; the corrosion rate of type 304 stainless steel increases rapidly above 95% concentration, while that of aluminum 3003 remains essentially constant to 100%. Nitric acid is a strong oxidizing agent and attacks most metals, such as iron, by oxidizing the metal to the oxide. A secondary effect of oxidation is the generation of hydrogen at the metal/acid interface, which can cause hydrogen embrittlement of some materials, for example, high-strength steels. Metals and alloys that are able to form adherent oxide films, such as austenitic stainless steels and aluminum alloys, are protected by their oxide films from corrosion by HNO3.
Corrosion by Hydrochloric Acid Hydrochloric acid is an important mineral acid with many uses, including acid pickling of steel, acid treatment of oil wells, chemical cleaning, and chemical processing. It is made by absorbing hydrogen chloride in water. Most acid is the by-product of chlorinations. Pure acid is produced by burning chlorine and hydrogen. Hydrochloric acid is available in technical, recovered, food-processing, and reagent grades. Reagent grade is normally 37.1%. Hydrochloric acid is a corrosive, hazardous liquid that reacts most metals to form explosive hydrogen gas and causes severe burns and irritation to the eyes and mucous membranes. Safe handling procedures are available from various manufacturers and from safety data sheets available from the Chemical Manufacturers Association (Arlington, VA).
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Zone
1
2 3 4 All zones (including 5)
Metal
ACI CN-7M(a)(b)(c) Monel(b)(c)(d) Copper(b)(c)(d) Nickel 200(b)(c)(d) Silicon bronze(b)(c)(d) Silicon cast iron(b)(e) Tungsten Titanium, grade 7 Titanium, grade 2(f) Silicon bronze(c)(d) Silicon cast iron(b)(e) Silicon cast iron(b)(e) Tungsten Titanium, grade 7(g) Platinum Tantalum Silver(b)(c) Zirconium(b)(c) Hastelloy B-2(b)(c) Molybdenum(b)(c)
(a) <2% HCl at 25 °C (75 °F). (b) No FeCl3 or CuCl2 contamination. (c) No free chlorine. (d) No aeration. (e) Contains chromium, molybdenum, and nickel. (f) <10% HCl at 25 °C (75 °F). (g) <5% HCl at boiling temperature
Fig. 23
Alloys with reported corrosion rates of <0.5 mm/year (<20 mils/year) in HCl
Materials Selection Guidelines for Hydrochloric Acid Concentrated HCl is transported and stored in rubber-lined tanks, although custom-fabricated polyester reinforced thermoset plastic (RTP) storage tanks have been used. Pipelines are usually plastic-lined steel. Processes involving aqueous acid are commonly carried out in glass-lined steel equipment. Nonmetallic materials are normally preferred because of the corrosive action of this strongest of acids on most metals. Candidate metals and alloys for handling HCl are shown in Fig. 23. Suitable materials are judged to be those with a corrosion rate under 0.5 mm/year (20 mils/year) when exposed to uncontaminated HCl. In practice, contamination is not uncommon and can be catastrophic. Selection of a candidate metal should be based on extensive corrosion testing or, preferably, field experience, using the grade of acid that will be available.
Corrosion by Hydrogen Fluoride and Hydrofluoric Acid Anhydrous hydrogen fluoride (AHF) and aqueous hydrofluoric acid (HF) are of great industrial importance. Anhydrous hydrogen fluoride is the foundation of the multibillion dollar fluorocarbon industry, which encompasses essentially all refrigerants, a fire-extinguishing agent, ultrasonic cleaning fluids, fluorocarbon plastics, and fluorocarbon elastomers. A popular process for alkylation of petroleum to enhance yields
Types of Corrosive Environments
of gasoline depends on the use of AHF. Aqueous HF is used in large quantities to pickle stainless steels, to acid treat wells, and to etch glass. Aqueous HF and AHF are hazardous chemicals. Fluoride salts, although added to potable waters to prevent tooth decay, are toxic in higher concentrations. Painful, persistent burns result from contact with aqueous HF or AHF, and inhalation of high concentrations of the vapors causes lung damage. The Occupational Safety and Health Administration has ruled that an employee’s exposure to HF vapor in any 8 h work shift of a 40 h work week shall not exceed a time-weighted average of 3 ppm HF vapor by volume.
Materials Selection Guidelines for Hydrofluoric Acid From the standpoint of corrosion, hydrofluoric acid is unique in its behavior compared to most other acids. The high-silicon cast irons, stoneware, and glass are generally resistant to most acids, but all are readily attacked by hydrofluoric acid. Magnesium shows rather poor corrosion resistance to many acids, but resists attack by hydrofluoric acid; in fact, shipping containers for this acid are often made of magnesium. Below 1% concentration, some attack occurs, but in concentrations of 5% and above, magnesium is practically immune to corrosion. The magnesium passivates, apparently because of the formation of a surface fluoride film, and corrosion of the metal is thus retarded. This same surface effect also applies to other metals and alloys that resist corrosion by hydrofluoric acid. Steels. Carbon steels have useful corrosion resistance from 64 to 100% HF at ambient temperatures. Low-alloy steels have higher corrosion rates than do plain carbon steels. Both hardened (>22 HRC) carbon and low-alloy steels are subject to hydrogen embrittlement in HF. Hydrogen blistering and stepwise cracking of plate and pipelines have been reported. Alloy 400 (Ni-31.5Cu-1.25Fe) possesses good corrosion resistance to all concentrations of HF up to temperatures of about 120 °C (250 °F). However, the presence of oxygen in the solution is detrimental and can cause stress-corrosion cracking (SCC). Oxidizing salts also tend to increase the corrosion of nickel-copper alloys. Copper-nickel alloy (70Cu-30Ni, C71500) is generally suitable for hot and cold dilute solutions and for higher concentrations up to about 65 °C (150 °F). Copper-nickel alloys are more sensitive to aeration and oxidizing salts in HF than are nickel-copper alloys. Lead shows fairly good resistance to HF in concentrations below 60% at room temperature. Stronger acids attack lead. High-Alloy Stainless Steels and High-Performance Nickel Alloys. Aerated HF or acid containing oxidizing salts, such as ferric or cupric ions, presents a problem, since the nickel-copper and copper-nickel
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alloys are unstable. In these cases, the use of nonmetallics and highalloy stainless steels such as 20Cb-3 should be considered. The Hastelloys (C-276, B-2, G, and G-3), Inconel 626, and Incoloy 825 also have good resistance to HF and to high-temperature HF vapors. Precious metals, including gold, silver, and platinum, are all more or less unaffected by HF of any concentration at ambient temperatures, and in most cases, to the boiling point or higher. Plastics, such as polyethylene, polypropylene, polyvinylidene fluoride and carbon-filled phenolics are limited to a maximum concentration of about 70% HF. Polyvinyl chloride is limited to about 50% concentration. Elastomers (rubbers) are also useful up to about 50 to 70% acid.
Corrosion by Phosphoric Acid Phosphoric acid (H3PO4) is obtained by two different means. The wet-process acid is obtained by reacting phosphate rock with concentrated sulfuric acid and concentrating the resulting dilute acid by evaporation. The furnace acid is obtained by calcining the phosphate rock to produce elemental phosphorus, which is then oxidized and reacted with water to produce phosphoric acid. This latter acid is very pure and is used as reagent grade. The wet-process acid is used to make phosphatic fertilizers and usually contains a number of impurities, such as HF, H2SO4, and SiO2. The percentage of these impurities depends on the source of the rock, the process of reaction with H2SO4, and the state of concentration of the H3PO4. Particularly sensitive to the impurities in wet-process acid are the stainless steels, which are widely employed in H3PO4 service. In addition, a given alloy can corrode at different rates in wet-process acids from different manufacturers. Because of these differences, any comparison between alloys in this type of acid must be made from tests conducted in the same batch of acid from the same source. The wetprocess acid can be considerably more corrosive than the reagent grade acid. When heated to temperatures above 200 °C (390 °F), phosphoric acid loses its water of constitution, and salts of the resulting dehydrated acids are used in preparing some liquid fertilizers and some detergents. Use of these salts in detergents, however, has been severely restricted in an attempt to reduce pollution by phosphates. Phosphoric acid is also used in soft drinks and flavored syrups, in pharmaceuticals, in water treatment, in animal feeds, and for pickling of rustproof metals.
Types of Corrosive Environments
Table 9
231
Results of plant corrosion test in a H3PO4 evaporator
Specimens were exposed for 42 days to 53% H3PO4 containing 1 to 2% H2SO4 and 1.2 to 1.5% flouride at 120 °C (250 °F). No pitting was observed. Corrosion rate Alloy
mm/yr
mils/yr
20Cb-3 C 825 Type 317 stainless steel 400 Type 316 stainless steel 600 B
0.12 0.13 0.16 0.26 0.64 1.1 >33(a) >33(a)
4.7 5 6.2 10.4 25 44 >1300(a) >1300(a)
(a) Specimen was destroyed
Materials Selection Guidelines for Phosphoric Acid Two of the most widely used materials for phosphoric acid service are type 316 stainless steel and the 20-type alloys (e.g., 20Cb-3). These alloys show very little attack in concentrations up to 85% and temperatures up to and including boiling. Other stainless steels used for wetprocess phosphoric acid service include 317L and iron-nickel-chromiummolybdenum alloys 904L, alloy 28, 20Mo-4, and 20Mo-6 (Chapter 6 provides chemical compositions). Nickel-base alloys such as Incoloy 825, Inconel 625, and Hastelloy alloys G-3 and C-276 are also used for severe service conditions. Table 9 compares the corrosion resistance of stainless steels and nickel-base alloys in dilute phosphoric acid. Lead and its alloys are also used at temperatures up to 200 °C (400 °F) and at concentrations up to 80% for pure and 85% for impure acid. Lead forms an insoluble phosphate on the surface that provides protection. High-silicon irons, glass, and stoneware show good resistance to pure acids. Rubber and plastics show good resistance but are largely limited to temperatures of about 80 °C (175 °F). Carbon and graphite are suitable at higher temperatures. Copper and high-copper alloys are not widely used in handling phosphoric acid. High-nickel-molybdenum alloys exhibit good resistance to pure acids but are attacked when aeration and oxidizing impurities are present. Aluminum, cast iron, steel, brass, and the ferritic and martensitic stainless steels do not possess sufficient corrosion resistance.
Corrosion by Organic Acids Organic acids constitute another group of important chemicals currently in use in industry. These acids are produced more as precursors for other chemicals than for end use as organic acids. Acetic acid is the
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most well-known member of the group and is produced in the largest volume; hence, it is emphasized in this section. The subject of corrosion by organic acids is complicated not only by the numerous acids to be considered but also by the fact that the acids typically are not handled as a pure chemical but, rather, as process mixtures with inorganic acids, organic solvents, salts, and mixtures of several organic acids.
Acetic Acid As stated above, acetic acid is the most important organic acid in terms of quantity produced and used. Many other organic acids show similar corrosion behavior and, in the absence of data, an approximation as to their corrosion behavior can usually be made by comparing them to acetic acid. The materials most suitable for acetic acid service are copper and the bronzes, 1100 and 3003 aluminum, type 304 and 316 stainless steel, 20-type alloys, high-silicon cast irons, Hastelloy C-276, and Chlorimet 3. These metals and alloys are specified for the majority of plant equipment handling acetic acid. Copper is widely used as a material of construction for acetic acid applications. More copper equipment is probably in service than any other material, although the austenitic stainless steels have replaced copper to some extent in recent years, particularly for oxidizing conditions. Copper-silicon alloys, copper-aluminum alloys, and the bronzes are also used where better mechanical or other properties are required. In general, copper and its alloys are suitable for concentrations below 99% with temperatures below boiling and for hot vapors, provided that aeration or oxidizing conditions are absent. Hot glacial acid (pure acetic acid containing less than 1% water) is destructive. In one interesting case involving corrosion tests in a column handling hot, 75% acetic acid vapors, copper showed practically no corrosion, and types 304 and 316 stainless steel were badly attacked. Copper presents a fabrication problem, because it is relatively difficult to obtain sound welds, and silversoldered joints often suffer galvanic corrosion at the expense of the copper. However, use of the inert-gas shielded tungsten-arc process for welding essentially eliminates both of these problems. This method of welding is now preferred over silver soldering for shop fabrication of copper equipment. Copper is also susceptible to erosion-corrosion. Aluminum has definite limitations. In general it is corroded by hot or warm acids that are not concentrated, but it shows good resistance to hot, glacial acetic acid provided that the acid does not contain substantial quantities of acetic anhydride. The corrosion resistance increases with concentration in the higher concentration range. Aluminum is suitable for all concentrations at room temperature. A small amount
Types of Corrosive Environments
(0.01%) of water must be present in order to develop a protective film. Increasing concentrations permit the use of higher temperatures. An advantage of aluminum equipment is that colored corrosion products are not formed, and thus there is less tendency for contamination of the product. High-Alloy Materials. The austenitic stainless steels and 20-type alloys are preferred over copper, nickel, and their alloys if aeration or oxidizing conditions exist. Type 304 is suitable for dilute and intermediate concentrations in boiling acid, but it shows appreciable attack in more concentrated boiling acid. Type 316 is suitable for boiling acid in all concentrations, although failures have occurred under reducing conditions. These 18-8 stainless alloys show borderline passivity in acetic acid and may exhibit either good resistance or destructive attack in the same corrosive environment. The presence of molybdenum in stainless steel greatly increases their corrosion resistance to acetic acid—even in the straight chromium steels, which in themselves are poorly resistant. Alloy 20 and Duriron (high-silicon cast iron) are resistant to all concentrations, including glacial, at up to and including boiling temperatures. Alloy 20 is considerably more resistant to acetic acid than types 304 and 316. It thus allows a greater margin of safety with regard to strongly reducing conditions, contamination of product, and impurities such as sulfuric or formic acids. Alloy 20 also possesses better resistance to erosion-corrosion than types 304 and 316, copper, and aluminum. Consequently, it is a desirable material for pumps and valves. Wrought alloy 20Cb-3 finds wide application in columns, bubble caps, pump shafts, tanks, and other vessels where corrosion conditions are too severe for the more common wrought materials. Duriron is widely used for pumping hot acetic acids. For temperatures above the atmospheric boiling point, Duriron pumps are sometimes kept warm by means of copper tubing that contains steam. Nickel and Monel are used if the temperature of the acid is below boiling and sometimes in glacial acid service. These alloys are susceptible to attack under oxidizing. Hastelloy C-276 and Chlorimet 3 are sometimes used for handling hot acetic acids when severely corrosive conditions exist and when the cost of these materials can be justified. Effect of Temperature. Acetic acid exhibits unusual corrosion behavior with regard to temperature effects. The corrosion resistance of metals and alloys usually decreases with increasing temperature. However, copper and the stainless steels switch position as the temperature is increased. For example, copper and 18-8 or 18-8Mo, stainless steels show low corrosion rates at room temperature. At boiling temperatures, the stainless steels are superior. However, at temperatures above boiling, the stainless steels are rapidly attacked, but copper shows negligible corrosion (provided that aeration is
233
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Corrosion: Understanding the Basics
absent). In fact, during an investigation of a corrosion problem involving a high-temperature, high-pressure heat exchanger it was found that one end of the exchanger was corrosive to 18-8Mo and not to copper; the opposite was true at the other end. This was due to the temperature drop along the tubes. The stainless alloys exhibit borderline passivity at high temperatures, and considerable care must be exercised in selection.
Other Organic Acids Formic acid shows corrosion characteristics similar to those of acetic acid. The main exception is aluminum, which should not be used in handling formic acid. Maleic and lactic acid are also similar to acetic acid, except that they are more aggressive toward sensitized 18-8 stainless and attack it intergranularly. Type 316 is a preferred material for handling fatty acids in the soap industry, particularly at higher temperatures.
Corrosion by Alkalis Common alkalis such as caustic soda (sodium hydroxide, NaOH) and potassium hydroxide are not particularly corrosive and can be handled in steel in most applications where contamination is not a problem. However, SCC and severe uniform corrosion must be guarded against at higher concentrations and temperatures. Sometimes rubber-based and other organic coatings and linings are applied to steel equipment to prevent iron contamination.
Materials Selection Guidelines for Alkalis Carbon steel is used to store up to 50% caustic soda at temperatures from ambient to about 95 °C (200 °F). Figure 24 shows ranges of behavior of steel as a function of temperature and caustic concentration. At lower temperatures in area A, carbon steel can be used without stress relief. At higher temperatures, carbon steel can be used with stress relief of welds and bends to lower the residual stresses and to decrease the likelihood of SCC. In area C at still higher temperatures, steel is no longer recommended; nickel alloys should be considered for these applications. The danger zone for steel in concentrated caustic can also be rationalized in terms of the potential-pH diagram. In Chapter 3, the potential-pH diagram for iron was presented, and a triangle of active corrosion was predicted in highly alkaline environments. If the conditions of the environment are within this region, either SCC or localized
150 300 140 280 Area C 125 260 115 240 105 220 95 200 180 80 70 160 60 140 Area B 120 50 40 100 Area A 25 80 60 15 No failure 5 40 Failure –5 20 0 –15 0 10 20 30 40 50 60 70 80 Concentration of NaOH, %
235
Temperature, °C
Temperature, °F
Types of Corrosive Environments
Fig. 24
NACE caustic soda chart superimposed over the data on which it is based (MTI Publication No. 15, 1985). Area A, carbon steel, no stress relief necessary; stress-relieved welded steam-traced lines. Area B, carbon steel; stress-relieved welds and bends. Area C, application of nickel alloys to be considered in this area
corrosion can occur. In severe cases, rapid uniform corrosion often referred to as caustic gouging can be observed. Nickel and nickel alloys are extensively used for combating corrosion by caustic soda. Nickel is suitable under practically all conditions of concentration and temperature. In fact, corrosion resistance to alkalis is almost directly proportional to the nickel content of an alloy. For example, as little as 2% Ni in cast iron is beneficial. Monel (alloy 400), the austenitic stainless steels containing 8 to 20% Ni, and other nickelbearing alloys are used in many applications involving high temperatures or control of contamination. Aluminum is unsuitable for equipment that handles caustic. Aluminum and its alloys are rapidly attacked by even dilute caustic environments.
References Selected References · Corrosion, Vol 13, ASM Handbook, ASM International, 1987, p 891–1370 · Corrosion Data Survey—Nonmetals Section, 5th ed., NACE International, 1975 · B. Craig and D. Anderson, Ed., Handbook of Corrosion Data, 2nd ed., ASM International, 1995 · J.R. Davis, Ed., Atmospheric Corrosion of Steels, Corrosion of Steels in Water, and Corrosion of Steels in Soils, ASM Specialty Handbook:
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· · · ·
Carbon and Alloy Steels, ASM International, 1996, p 393–407, 408–429, and 430–438 M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill Book Company, 1986 D.L. Graver, Ed., Corrosion Data Survey—Metals Section, 6th ed., NACE International, 1985 G. Kobrin, A Practical Manual on Microbiologically Influenced Corrosion, NACE International, 1993 L.S. Van Delinder, Ed., Corrosion Basics—An Introduction, NACE International, 1984
Corrosion: Understanding the Basics J.R. Davis, editor, p237-300 DOI: 10.1361/cutb2000p237
CHAPTER
Copyright © 2000 ASM International® All rights reserved. www.asminternational.org
6
Corrosion Characteristics of Structural Materials ALL MATERIALS, including metals and their alloys, plastics, and ceramics, are susceptible to corrosion or some form of environmental degradation. Although no single material is suitable for all applications, usually there are a variety of materials that will perform satisfactorily in a given environment. As described in Chapter 8, selection of the proper material may not simply be a matter of selecting the most corrosion resistant material. For example, carbon steels are the most common materials of construction, not because of their corrosion resistance, which is usually fair to poor, but because of their excellent mechanical properties, weldability, and low cost. When used in corrosive environments, steels usually require some form of corrosion protection, such as protective coatings, cathodic protection, or a combination of these corrosion prevention methods. The intent of this chapter is to (a) review the corrosion behavior of the major classes of metals and alloys as well as some nonmetallic materials, (b) describe typical corrosion applications, (c) present some unique weaknesses of various types of materials, (d) point out some unique material characteristics that may be important in material selection, and (e) discuss, where appropriate, the characteristic forms of corrosion that attack specific materials. Complementary information on the forms of corrosion that affect different alloys can be found in Chapter 4. Two excellent references for more detailed information
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on the corrosion behavior of various alloys are Corrosion, Volume 13 of the ASM Handbook, and the Handbook of Corrosion Data, 2nd edition, both of which are published by ASM International. The selected references at the conclusion of this chapter should also be consulted.
Carbon Steels Carbon steels contain up to approximately 1.0% C and alloy contents of generally less than 2% by weight. Despite its relatively limited corrosion resistance, very large amounts of carbon steels are used in marine applications, nuclear and fossil fuel power plants, and transportation, chemical processing, mining, construction, and metal- processing equipment. All of these areas present unique corrosion problems.
Corrosive Service As a rule, only the low-carbon, or mild, steels containing 0.08 to 0.28% C are considered for resistance to corrosion. They are generally more corrosion resistant than the medium-carbon (0.28 to 0.55% C) and high-carbon (0.50 to 1.0% C) groups. More importantly, they are more amenable to welding and forming, a common requirement for building structures of a variety of types. Without some sort of surface protection, carbon steels are hardly worth considering for resisting attack by very aggressive chemicals, simply because this kind of attack is so very rapid. However, carbon steel in vast quantities has been successfully used in corrosive atmospheres and waters. The two conditions necessary to initiate corrosion of low-carbon steel within natural environments are water and oxygen. After these essentials, a number of variables can affect the corrosion process. For example, samples of mild steels that are wholly immersed corrode faster if uninhibited water is moving around them than if the water is stagnant, but less rapidly where velocity is high if an inhibitor is present. Another of the many variables is the process called “cycling.” Water pipes and tanks corrode much more slowly if the immersed surface remains completely submerged as opposed to alternating between submersion and partial or total exposure. Carbon steels perform well in dry, rural atmospheres, but the rate of corrosion increases quickly in high-humidity saline or industrial atmospheres. The useful service life of carbon steels has been recorded for boiler service of up to 25 years where the conditions are controlled. It is thus obvious that stainless steels, copper-base alloys, and other highly
Corrosion Characteristics of Structural Materials
corrosion-resistant alloys are not always required for many such applications. Corrosive Environments. There are a number of major corrosive conditions into which carbon steels can be successfully introduced. These include the following: · Atmospheric corrosion, including humidity, and both natural and man-made pollutants · Soil corrosion, as determined by such factors as moisture content, level of electrical conductivity, acidity level, amount of dissolved salts, and aeration · Corrosion in concrete, as caused primarily by chloride ions, and most successfully combated by cathodic protection · Boiler service, which is a specialized form of aqueous corrosion that also involves elevated temperatures · Corrosion in molten nitrate salts, which are used for heat treatment baths (plain carbon and low-alloy steels form protective iron oxide films that effectively protect the metal surface to approximately 500 °
